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A COLLEGE TEXT-BOOK 



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OF 



CHEMISTRY 



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BY 

IRA REMSEN 

President of the Johns Hopkins University 



t 




NEW YORK 

HENRY HOLT AND COMPANY 
1901 



THE LIBRARY OF 

CONGRESS, 
Two Copies Received 

OCT. 2 1901 

Copyright entry 
CLASS 0>xXc No. 

I%I7Z- 
oopy a . 



m. 



Copyright, 1901, 

BT 

HENRY HOLT & CO. 



ROBERT DRUMMOND, PRINTER, NEW YORK 



/97 



PREFACE. 

This book is intended to fill a place between "An In- 
troduction to the Study of Chemistry " and the " Inorganic 
Chemistry " of this series. It is distinctly more advanced 
than the former and distinctly less advanced than the 
latter. The "Inorganic Chemistry" has been largely 
used in its preparation, but "every part has been carefully 
revised. Especial attention has been given to the parts 
dealing with matters pertaining to physical chemistry, and 
it is hoped that the treatment will commend itself to those 
competent to judge. 

In the opinion of the author the time has not yet come 
for the abandonment of the study of elements and their 
compounds in what some are pleased to call the old- 
fashioned way. Indeed it seems essential that such study 
must always form the basis of the higher or spiritual study 
of chemistry. All chemists are thankful for the new con- 
ceptions that have been given to them in the last few years, 
and students must be made familiar with them in a 
general way. 

Ira Remsen. 

Baltimore, September 11, 1901. 



67 



CONTENTS. 



CHAPTER I. 

CHEMICAL AND PHYSICAL CHANGE— EARLIEST CHEMICAL KNOWLEDGE 
LAW OP THE INDESTRUCTIBILITY OF MATTER — LAW OF DEFI- 
NITE PROPORTIONS — LAW OF MULTIPLE PROPORTIONS— THE 

ELEMENTS. 

PAGF 

Matter and Energy— Chemical Change— Physical Change — 
Physics and Chemistry — Earliest Chemical Knowledge — 
Law of the Indestructibility of Matter — Conservation of 
Energy — Composition of Matter — Elements — Chemical Ac- 
tion — Chemical Affinity — Chemical Compounds and Mechan- 
ical Mixtures — Qualitative and Quantitative Study of Chem- 
ical Changes — Law of Definite Proportions — Law of Multiple 
Proportions — Combining Weights of the Elements — The 
Elements, their Symbols and Atomic Weights — Symbols of 
Compounds — Chemical Equations — The Scope of Chemistry 
— Chemical Action accompanied by other Kinds of Action. 

Experiments : Chemical Change caused by Heat — Chemical 
Changes can be effected by an Electric Current — Mechanical 
Mixtures and Chemical Compounds— Other Examples of 
Chemical Action 1 

CHAPTER II. 

A STUDY OF THE ELEMENT OXYGEN. 

Historical — Occurrence — Preparation — Physical Properties — 
Chemical Properties — Burning in the Air and Burning in 
Oxygen — Combustion — Kindling Temperature — Slow Oxida- 
tion — Heat of Combustion — Heat of Decomposition — Chemi- 
cal Energy and Chemical Work — Oxides. 

Experiments : Preparation of Oxygen — Measurement of the 
Volume of Gases — Determination of the Amount of Oxygen 
liberated wheu a known Weight of Potassium Chlorate is 
decomposed by Heat — Physical Properties of Oxygen — 



vi CONTENTS. 

PAGE 

Chemical Properties of Oxygen — Oxygen is used up in Com- 
bustion — The Products of Combustion weigh more than the 
Body burned, 32 

CHAPTER III. 

A STUDY OF THE ELEMENT HYDROGEN. 

Historical — Occurrence — Preparation — Physical Properties — 
Chemical Properties — Comparison of Oxygen and Hydrogen. 

Experiments : Preparation of Hydrogen — Something besides 
Hydrogen is formed — Determination of the Amount of 
Hydrogen evolved when a Known Weight of Zinc is dis- 
solved in Sulphuric Acid — Hydrogen is purified by passing 
through a Solution of Potassium Permanganate — Hydrogen 
passes readily through Porous Vessels — Diffusion — Chemical 
Properties of Hydrogen — Product formed when Hydrogen is 
burned — Reduction, 54 

CHAPTER IV. 

STUDY OF THE ACTION OF HYDROGEN ON OXYGEN. 

Burning of Hydrogen — Method of Dumas — Morley's Method — 
Eudiometric Method — Calculation of the Result obtained in 
exploding Mixtures of Hydrogen and Oxygen — Determina- 
tion of the Volume of Water Vapor formed by Union of 
Definite Volumes of Hydrogen and Oxygen — Heat evolved in 
the Union of Hydrogen and Oxygen — Applications of the 
Heat formed by the Combination of Hydrogen and Oxygen — 
Oxyhydrogen Light — Summary. 

Experiments: Composition of Water — Eudiometric Experi- 
ments — Oxyhydrogen Blow-pipe, 72 

CHAPTER V. 

WATER. 

Historical— Occurrence — Formation of Water and Proofs of its 
Composition — Properties of Water — Chemical Properties of 
Water — Water as a Solvent— Solution as an Aid to Chemi- 
cal Action— Natural Waters— What constitutes a Bad Drink- 
ing Water — Purification of Water. 

Experiments: Organic Substances contain Water — Water of 
Crystallization — Efflorescent Salts — Deliquescent Salts — 
Purification of Water by Distillation, 83 



CONTENTS. vii 

CHAPTER VI. 

CONSTITUTION OF MATTER — ATOMIC THEORY — ATOMS AND MOLE- 
CULES—CONSTITUTION — VALENCE. 

PAGE 

The Atomic Theory as proposed by Dalton — Use and Value of a 
Theory — Atomic Weights and Combining Weights — Mole- 
cules — Avogadro's Hypothesis — Distinction between Mole- 
cules and Atoms — Molecular Weights — Deduction of Atomic 
Weights from Molecular Weights — Exact Atomic Weights 
determined by the Aid of Analysis — Molecular Formulas — 
Constitution — Valence — Replacing Power of Elements. 

Experiments: Method of Dumas — Method of Victor Meyer, . 97 

CHAPTER VII. 

OZONE — ALLOTROPY — NASCENT STATE — HYDROGEN DIOXIDE. 

Occurrence — Preparation — Properties — Relation between Oxygen 
and Ozone — Ozone in the Air — Allotropy — Varying Number 
of Atoms in the Molecules of one and the same Element — 
Nascent State — Hydrogen Dioxide or Hydrogen Peroxide — 
Properties — Occurrence in the Air — Characteristic Reactions. 

Experiments : Ozone — Hydrogen Dioxide, 116 

CHAPTER VIII. 

CHLORINE — HYDROCHLORIC ACID. 

Historical — Occurrence of Chlorine — Preparation — Deacon's Pro- 
cess — Laboratory Process — Weldon's Process — Electrolytic 
Process — Properties — Different Kinds of Action — Chlorine 
Hydrate and Liquid Chlorine — Application of Chlorine — Hy- 
drochloric Acid — Historical — Study of the Action of Hydro- 
gen upon Chlorine — Preparation — Properties — Chemical 
Action of Hydrochloric Acid. 

Experiments: Preparation of Chlorine — Chlorine decomposes 
Water in the Sunlight — Chlorine Hydrate — Formation of 
Hydrochloric Acid— Preparation of Hydrochloric Acid, . . 126 

CHAPTER IX. 

COMPOUNDS OF CHLORINE WITH OXYGEN AND WITH HYDROGEN 
AND OXYGEN. 

General — Principal Reactions for Making Compounds of Chlorine 
with Hydrogen and Oxygen — Hypochlorous Acid — Perchloric 
Acid — General — Compounds of Chl'Tine with Oxygen. 



i 



vni CONTENTS. 



PAGE 

Experiments : Chloric Acid and Potassium Chlorate — Perchloric 

Acid, 14 6 



CHAPTER X. 

ACIDS — BASES — NEUTRALIZATION — SALTS. 

General— A Study of the Act of Neutralization — General State- 
ments— Definitions— Distinction between Acids and Bases — 
Metals or Base- forming Elements — Constitution of Acids and 
Bases — Constitution of Salts— Basicity of Acids — Acidity of 
Bases — Salts — Acid Properties and Oxygen — Nomenclature 
of Acids — Nomenclature of Bases — Nomenclature of Salts. 

Experiments: Neutralization of Acids and Bases; Formation of 

Salts — Study of the Products formed, 153 



CHAPTER XI. 

NATURAL CLASSIFICATION OF THE ELEMENTS — THE PERIODIC LAW. 

Historical — Arrangement of the Elements — Connection between 
the Position of the Elements in the. Natural System and 
their Chemical Properties — Plan to be followed, .... 173 

CHAPTER XII. 

THE ELEMENTS OF FAMILY VII, GROUP B: 
FLUORINE — CHLORINE — BROMINE— IODINE. 

General — Bromine — Occurrence — Preparation — Properties — 
Chemical Conduct of Bromine — Uses of Bromine — Hydro- 
bromic Acid— Properties— Compounds of Bromine with Hy- 
drogen and Oxygen — Iodine — Occurrence — Preparation — 
Properties — Hydriodic Acid — Properties — Iodic Acid— Iodine 
Pentoxide or Iodic Anhydride — Anhydrides, or Acidic Oxides 
— Periodic Acid — Compounds of Iodine with Chlorine — 
Compound of Iodine with Bromine — Fluorine — Occurrence 
— Properties — Hydrofluoric Acid — Family VII, Group A — 
Manganese. 

Experiments : Preparation of Bromine — Hydrobromic Acid — 
Iodine — Iodine can be detected by -Means of its Action upon 
Starch-paste — Iodic Acid — Hydrofluoric Acid, 186 



CONTENTS. ix 



CHAPTER XIII. 

THE ELEMENTS OF FAMILY Vr, GROUP B : 
SULPU tJIl — SELENIUM — TELLURIUM. 

PAGE 

Introductory — Sulphur — Occurrence — Extraction of Sulphur 
from its Ores — Properties — Uses of Sulphur — Hydrpgen Sul- 
phide, Sulphuretted Hydrogen — Properties — Action of Hy- 
drogen Sulphide upon Solutions of Salts, Use in Chemical 
Analysis — Hydrosulphides — Hydrogen Persulphide — Com- 
pounds of Sulphur with Members of the Chlorine Group — 
Selenium — Occurrence — Properties — Tellurium— Occurrence 
— Properties. 

Experiments: Properties of Sulphur — Hydrogen Sulphide, . 205 



CHAPTER XIV. 

COMPOUNDS OF SULPHUR, SELENIUM, AND TELLURIUM WITH OXYGEN 
AND WITH OXYGEN AND HYDROGEN. 

Introductory — Sulphuric Acid — Pure Sulphuric Acid — Tetrahy- 
droxyl Sulphuric Acid — Normal Sulphuric Acid — Disulphuric 
Acid, Pyrosulphuric Acid — Sulphurous Acid — Hyposulphur- 
ous Acid — Thiosulphuric Acid — Other Acids of Sulphur — 
Persulphuric Acid — Caro's Reagent — Compounds of Sulphur 
with Oxygen — Sulphur Dioxide — Sulphur Trioxide — Sul- 
phuryl Chloride — Chlorsulphuric Acid, or Sulphuryl- 
hydroxyl Chloride — Compounds of Selenium and Tellurium 
with Oxygen and with Oxygen and Hydrogen — Selenious 
Acid — Selenic Acid — Selenium Dioxide— Tellurious Acid — 
Telluric Acid — Oxides of Tellurium — Family VI, Group A. 

Experiments : Manufacture of Sulphuric Acid — Sulphurous 
Acid and Sulphur Dioxide — Sulphurous Acid is a Reducing 
Agent— Sulphur Trioxide 222 



CHAPTER XV. 

NITROGEN— THE AIR — ARGON, ETC. 

Nitrogen — General — Occurrence of Nitrogen — Preparation — 
Properties — The Air — Analysis of Air — Air and Life — Pure 
and Impure Air — Liquid Air— Argon — Other Gases in Air. 

Experiments : Preparation of Nitrogen — Analysis of Air, . . 250 



CONTENTS. 



CHAPTER XVI. 

COMPOUNDS OF NITROGEN WITH HYDROGEN — WITH HYDROGEN AND 
OXYGEN— WITH OXYGEN, ETC. 

PAGE 

General Conditions which, give Rise to the Formation of the Sim- 
pler Compounds of Nitrogen — Relations between the Princi- 
pal Compounds of Nitrogen — Ammonia — Properties — Com- 
position of Ammonia — Ammonium Amalgam — Hydrazine — 
Hydroxy lamine — Triazoic Acid — Nitric Acid — Red Fuming 
Nitric Acid — Nitrous Acid — Hyponitrous Acid — Nitrous 
Oxide — Nitric Oxide — Nitrogen Trioxide — Nitrogen Per- 
oxide — Nitrogen Pentoxide — Compounds of Nitrogen with 
the Elements of the Chlorine Group. 

Experiments : Preparation and Properties of Ammonia — Am- 
monia burns in Oxygen. — Ammonia forms Ammonium Salts 
with Acids — Composition of Ammonia — Preparation and 
Properties of Nitric Acid — Nitric Acid gives up Oxygen 
readily, and is hence a Good Oxidizing Agent — Metals dis- 
solve in Nitric Acid, forming Nitrates — Nitrates are decom- 
posed by Heat — Nitrates are soluble in "Water — Nitric Acid 
is reduced to Ammonia by Nascent Hydrogen — Nitrous Acid 
— Nitrous Oxide — Nitric Oxide — Nitrogen Trioxide — Nitro- 
gen Peroxide, 267 



CHAPTER XVII. 

ELEMENTS OP FAMILY V, GROUP B: 

PHOSPHORUS — ARSENIC— ANTIMONY— BISMUTH— THE ELEMENTS AND 

THEIR COMPOUNDS WITH HYDROGEN. 

General — Phosphorus — Occurence — Preparation — Properties — 
Applications of Phosphorus — Compounds of Phosphorus with 
Hydrogen — Phosphine, Gaseous Phosphuretted Hydrogen — 
Arsenic — Occurrence — Preparation — Properties — Arsine, 
Arseniuretted Hydrogen — Antimony — Occurrence — Proper- 
ties — Applications of Antimony — Stibine — Methods of Dis- 
tinguishing between Arsenic and Antimony — Bismuth — Oc- 
currence — Compounds of the Members of the Phosphorus 
Group with the Members of the Chlorine Group — Phos- 
phorus Trichloride — Phosphorus Pentachloride — Arsenic 
Trichloride — Compounds of Antimony and Chlorine — Bis- 
muth and Chlorine. 



CONTENTS. xi 

PAGE 

EXPERIMENTS : Phosphorus — Phosphorus abstracts Oxygen from 
other Substances — Phosphine — Arsenic — Arsine — Marsh's 
Test for Arsenic — Antimony — Stibine — Bismuth — Phos- 
phorus Trichloride — Phosphorus Pentachloride, .... 298 



CHAPTER XVIII. 

COMPOUNDS OF THE ELEMENTS OF THE PHOSPHORUS GROUP WITH 
OXYGEN AND WITH OXYGEN AND HYDROGEN. 

Introduction — Phosphoric Acid, Orthophosphoric Acid — Prop- 
erties — Pyrophosphoric Acid — Metaphosphoric Acid — Phos- 
.phorous Acid — Hypophosphoric Acid — Hypophosphorous 
Acid — Phosphorus Pentoxide, Phosphoric Anhydride — Phos- 
phorus Trioxide or Phosphorous Anhydride — Phosphorus 
Oxychloride — Arsenic Acid — Arsenious Acid — Arsenic Tri- 
oxide — Arsenic Pentoxide — Sulphides — Arsenic Disulphide 
— Arsenic Trisulphide — Arsenic Pentasulphide — Antimonic 
Acid — Antimony Trioxide — Salts of Antimony — Antimony 
Tetroxide — Antimony Pentoxide — Antimony Trisulphide — 
Antimony Pentasulphide — Oxychlorides of Antimony — 
Oxides of Bismuth — Salts of Bismuth — Bismuth Dioxide — 
Bismuth Pentoxide — Bismuth Trisulphide — Bismuth Oxy- 
chloride — Family V, Group A — Vanadium — Vanadic Acid — 
Columbium — Tantalum — Didymium — Boron — General — Oc- 
currence — Preparation — Properties — Boron Trichloride — 
Boric Acid — Salts of Boron — Nitrogen Boride. 

Experiments: Phosphoric Acid— Arsenic Acid — Reduction of 
Arsenic Trioxide — Sulphides of Arsenic — Sulphides of Anti- 
mony — Oxychlorides of Antimony — Basic Nitrates of Bis- 
muth—Boron, 320 



CHAPTER XIX. 

CARBON AND ITS SIMPLER COMPOUNDS WITH HYDROGEN AND 
CHLORINE. 

Introductory — Occurrence of Carbon — Diamond — Graphite — 
Amorphous Carbon — Coal — Diamond, Graphite, and Char- 
coal are Different Forms of the Element Carbon — Chem- 
ical Conduct of Carbon — Compounds of Carbon with Hydro- 
gen, or Hydrocarbons. 

Experiments : Carbon — Bone-black Filters — Charcoal absorbs 



xii CONTENTS. 

PAGE 

Gases— Carbon combines with Oxygen to form Carbon Diox- 
ide—Carbon reduces some Oxides when heated with them — 
Hydrocarbons, 345 



CHAPTER XX. 

SIMPLER COMPOUNDS OF CARBON WITH OXYGEN, AND WITH OXY- 
GEN AND HYDROGEN. 

' General — Carbon Dioxide — Preparation — Properties — Relations 
of Carbon Dioxide to Chemical Energy — Respiration— Car- 
bon Dioxide and Life — Energy stored up in Plants — Car- 
bonic Acid and Carbonates — Carbon Monoxide. 

Experiments : Carbon Dioxide is formed when a Carbonate is 
treated with an Acid — Preparation and Properties of Carbon 
Dioxide — Carbon Dioxide is given off from the Lungs— For- 
mation of Carbonates — Preparation and Properties of Carbon 
Monoxide — Carbon Monoxide is a Good Reducing Agent, . 359 



CHAPTER XXI. 

ILLUMINATION — FLAME — BLOWPIPE. 
COMPOUNDS OF CARBON WITH NITROGEN AND SULPHUR. 

Introduction — Illuminating Gas, Coal-gas — Flames — Kindling 
Temperature of Gases — Miner's Safety Lamp — Structure of 
Flames— Blowpipe — Causes of the Luminosity of Flames — 
Bunsen Burner — Compounds of Carbon with Nitrogen and 
with Sulphur — Cyanogen — Hydrocyanic Acid, Prussic Acid 
— Cyanic Acid — Carbon Disulphide — Sulphocyanic Acid — 
Constitution of Cyanogen and its Simpler Compounds. 

Experiments : Coal-gas — Oxygen barns in an Atmosphere of a 
Combustible Gas — Kindling Temperature of Gases — The 
Blowpipe and its Uses — Cyanogen, 372 



CHAPTER XXII. 

ELEMENTS OF FAMILY IV, GROUP A : 
SILICON — TITANIUM — ZIRCONIUM— CERIUM— THORIUM. 

General — Silicon — Occurrence — Preparation — Silicon Hydride — 
Titanium — Zirconium — Thorium — Cerium — Compounds of 
the Elements of the Silicon Group with those of the Chlo- 



CONTENTS. xin 

PAGE 

rine Group — Silicon Tetrachloride — Silicon Hexacbloride — 
Silicon Tetrafluoride — Compounds of the Members of the 
Silicon Group with Oxygen, and with Oxygen and Hydrogen 
— Silicon Dioxide — Properties — Uses — Silicic Acid— Poly- 
silicic Acids — Disilicic Acid — Trisilicic Acids — Titanium Di- 
oxide— Silicides — Family IV, Group B. 
Experiments : Silicon — Silicon Tetrafluoride and Fluosilicic 
Acid — Silicic Acid, 387 



CHAPTER XXIII. 

CHEMICAL ACTION. 

Retrospective — Classification of Reactions of the Elements 
and Compounds Studied — Kinds of Chemical Reactions — Di- 
rect Combination — Direct Decomposition — Metathesis — The 
Cause of Chemical Reactions — An Ideal Chemical Reaction 
— Influence of Mass — Reactions may be Complete if One of 
the Products formed is Insoluble or Volatile — Dissociation — 
Electrolysis— Electrolytic Dissociation — Definition of Acids 
and Bases in Terms of the Theory of Electrolytic Disso- 
ciation — Raoult's Methods for the Determination of Molec- 
ular Weights — Determination of the Extent of Dissociation 
of a Dissolved Substance — Osmotic Pressure — Relations be- 
tween Specific Heat and Atomic Weights —Exceptions to the 
Law of Specific Heats, 402 

CHAPTER XXIV. 

BASE-FORMING ELEMENTS — GENERAL CONSIDERATIONS. 

Introductory — Metallic Properties — Order in which the Base- 
forming Elements will be taken up — Occurrence of the 
Metals — Extraction of the Metals from their Ores — The 
Properties of the Metals — Compounds of the Metals — Chlo- 
rides — Formation of Salts in General — General Properties of 
the Chlorides — The so-called Double Chlorides and similar 
Compounds of Fluorine, Bromine, and Iodine — Different 
Chlorides of the same Metal — Oxides — Different Oxides of 
the same Metal — Hydroxides— Decomposition of Salts by 
Bases — Sulphides — Hydrosulphides — Sulpho-salts — Nitrates 
— Chlorates — Sulphates — Carbonates — Phosphates — Sili- 
cates. 



xiv CONTENTS. 

PAGE 

Experiments : Chlorides, Bromides, and Iodides — Hydroxides 
— Sulphates — Reduction of Sulphates to Sulphides— Carbon- 
ates, 457 



CHAPTER XXV. 

ELEIJENTS OF FAMILY 1, GROUP A: 

THE ALKALI METALS :— LITHIUM — SODIUM— POTASSIUM — RUBIDIUM 

— CAESIUM — AMMONIUM. 

General — Potassium — Occurrence — Preparation — Properties — 
Potassium Hydride — Potassium Fluoride, Chloride,Bromide, 
Iodide — Properties — Applications — Potassium Hydroxide — 
Potassium Oxide — Potassium Hydrosulphide — Potassium Ni- 
trate — Applications — Gunpowder — Potassium Nitrite— Po- 
tassium Chlorate — Potassium Perchlorate — Potassium Cyan- 
ide — Potassium Sulphocyanate — Potassium Sulphate — 
Primary, or Acid, Potassium Sulphate — Sulphites — Carbon- 
ates — Acid Potassium Carbonate — Phosphates — Potassium 
Silicate — Rubidium — Caesium — Sodium — Occurrence — Prep- 
aration — Properties — Applications — Sodium Hydride — So- 
dium Chloride — Sodium Hydroxide — Oxides — Sodium Per- 
oxide — Sodium Sulphantimonate — Sodium Nitrate — Sodium 
Sulphate— Sodium Thiosulphate — Sodium Carbonate — Prop- 
erties—Applications—The Le Blanc Process for the Manufac- 
ture of Sodium Carbonate — Ammonia Process for the Manu- 
facture of Soda— Mono-Sodium Carbonate, Primary Sodium 
Carbonate— Phosphates— Sodium Borate — Sodium Silicate 
— Lithium — Ammonium Salts — Ammonium Chloride — Am- 
monium Sulphide — Ammonium Nitrate — Sodium-ammonium 
Phosphate— Reactions of the Members of the Sodium Group 
which are of Value in Chemical Analysis — Flame Reactions 
and the Spectroscope. 

Experiments : Potassium Salts— Sodium Salts, 456 



CHAPTER XXVI. 

ELEMENTS OF FAMILY II, GROUP A : 
GLTJCINUM — MAGNESIUM - CALCIUM — STRONTIUM— BARIUM [ERBIUM]. 

General — Calcium Sub-group — Calcium — Occurrence — Prepara- 
tion — Properties — Calcium Chloride — Calcium Fluoride — 
Calcium Oxide — Calcium Hydroxide — Applications — Bleach- 



CONTENTS. XV 

PAGB 

ing-powder — Applications — Calcium Carbonate — Applica- 
tions — Calcium Sulphate — Applications — Calcium Phos- 
phates — Calcium Silicate — Glass — Mortar — Calcium Sul- 
phide — Calcium Nitride — Calcium Carbide — Strontium — 
Occurrence and Preparation — Properties — Compounds of 
Strontium — Barium — Occurrence and Preparation — Proper- 
ties — Barium Chloride — Barium Hydroxide — Barium Oxide 
— Barium Peroxide or Dioxide — Barium Sulphide — Barium 
Sulphate — Barium Carbonate — Phosphates of Barium — Re- 
actions which are of Special Value in Analysis — Magnesium 
Sub-group — Glucinum — Occurrence and Preparation — Mag- 
nesium — Occurrence — Preparation — Properties — Applica- 
tions — Compounds of Magnesium — Magnesium Chloride — 
Magnesium Oxide — Magnesium Sulphate — Applications — 
Magnesium Carbonate — Phosphates — Borates — Silicates — 
Magnesium Silicide — Reactions of Magnesium Salts which, 
are of Special Value in Chemical Analysis — Erbium. 
Experiments : Calcium Salts — Magnesium and its Salts, . . 492 

CHAPTER XXVII. 

ELEMENTS OF FAMILY III, GROUP A : 

ALUMINIUM — SCANDIUM — YTTRIUM — YTTERBIUM— SAMARIUM — 

HELIUM. 

General — Aluminium — Occurrence — Preparation — Properties — 
Applications — Aluminium Chloride — Aluminium Hydrox- 
ide — Aluminates — Aluminium Oxide — Aluminium Sulphate 
— Alums— Potassium Alum, Potassium-aluminium Sulphate 
— Applications — Ammonium Alum, Ammonium-aluminium 
Sulphate — Sodium Alum — Aluminium Silicates — Alloys- 
Reactions of Aluminium Salts which are of Special Value in 
Chemical Analysis — Other Members of Family III, Group 
A — Scandium — Yttrium — Ytterbium — Samarium — Helium. 

Experiments : Aluminium Chloride, 521 



CHAPTER XXVIII. 

ELEMENTS OF FAMILY I, GROUP B : 
COPPER — SILVER— GOLD. 

General— Copper — General— Forms in which Copper Occurs in 
Nature — Metallurgy of Copper — Properties — Applications — 
Alloys — Cuprous Chloride — Cupric Chloride— Cuprous Hy- 



xvi CONTENTS. 



droxide — Cuprous Oxide — Cupric Hydroxide — Cupric Oxide 
— Cupric Sulphate — Applications — Cupric Nitrate — Cupric 
Arsenite — Cupric Carbonates — Cuprous Sulphide — Cupric 
Sulphide — Copper-plating — Reactions which are of Special 
Value in Chemical Analysis — Silver — General— Forms in 
which Silver occurs in Nature — Metallurgy of Silver — Puri- 
fication — Properties — Allotropic Forms of Silver — Alloys of 
Silver — Argentous- Chloride — Silver Chloride, Argentic 
Chloride — Silver Bromide and Iodide — Application of the 
Chloride, Bromide, and Iodide of Silver in the Art of Pho- 
tography — Silver Triazoate — Silver Oxide — Other Oxides of 
Silver — Sulphides of Silver— Silver Nitrate, Argentic Ni- 
trate — Silver Cyanide — Reactions which are of Special Value 
in Chemical Analysis — Gold — General — Forms in which Gold 
occurs in Nature — Metallurgy of Gold — Properties — Alloys 
of Gold— Chlorides of Gold— Gold Sulphide. 
Experiments : Copper and its Salts— Silver and its Salts, . . 537 



CHAPTER XXIX. 

ELEMENTS OF FAMILY II, GROUP B: 
ZINC — CADMIUM — MERCURY. 

General — Zinc — General — Forms in which it occurs in Nature — 
Metallurgy — Properties — Applications — Allay*— Zinc Chlo- 
ride — Zinc Hydroxide — Zinc Oxide — Zinc Sulphide — Zinc 
Sulphate — Zinc Carbonate — Reactions which are of Special 
Value in Chemical Analysis — Cadmium — General — Prepa- 
ration and Properties — Cadmium Sulphide — Analytical Reac- 
tions — Mercury — General — Forms in which Mercury occurs 
in Nature — Metallurgy of Mercury — Properties — Applica- 
tions — Amalgams — Mercurous Chloride — Mercuric Chloride, 
or Corrosive Sublimate — Mercurous Iodide — Mercuric Iodide 
— Mercuric Oxide — Mercuric Sulphide — Mercuric Cyanide — 
Mercurous Nitrate — Mercuric Nitrate — Compounds formed 
by Salts of Mercury with Ammonia — Reactions which are of 
Special Value in Chemical Analysis. 

ELEMENTS OF FAMILY III, GROUP B : 
GALLIUM — 1NDI UM — THALLIUM. 

General — Gallium — Indium — Thallium. 

Experiments : Zinc and its Salts — Mercury and its Salts, . . 559 



CONTENTS. xv ll 



CHAPTER XXX. 



ELEMENTS OF FAMILY IV, GROUP B : 
GERMANIUM — TIN— LEAD. 

PAGE 

General — Germanium — Tin — General — Occurrence — Metallurgy 
— Properties — Applications — Alloys — Stannous Chloride — 
Stannic Chloride — Stannic Hydroxide — Stannic Oxide — 
Stannous Sulphide — Stannic Sulphide — Reactions which are 
of Special Value in Chemical Analysis — Lead — General — 
Forms in which Lead occurs in Mature — Metallurgy — Prop- 
erties — Applications — Lead Chloride — Lead Tetrachloride — 
Lead Iodide — Oxides of Lead — Lead Suboxide — Lead Oxide 
— Lead Peroxide — Storage Battery — Lead Sulphide — Lead 
Nitrate — Lead Carbonate — Lead Sulphate — Reactions which 
are of Special Value in Chemical Analysis — Lanthanum — 
Cerium — Didymium — Praseodymium and Neodymium. 

Experiments : Tin and its Compounds— Lead and its Com- 
pounds, 576 



CHAPTER XXXI. 

ELEMENTS OF FAMILY VI, GROUP A : 
CHROMIUM — MOLYBDENUM— TUNGSTEN — URANIUM. 

General — Chromium — General — Forms in which Chromium oc- 
curs in Nature — Preparation — Properties — Chromic Chloride 
— Chromic Hydroxide — Chromic Oxide — Chromic Sulphate — 
Chrome-Alums — Chromic Acid and the Chromates — Potas- 
sium Chromate — Potassium Bichromate — Chromium Trioxide 
— Relations between the Chromates and Bichromates — Sodium 
Chromate and Sodium Bichromate — Barium Chromate — Lead 
Chromate — Reactions which are of Special Value in Chem- 
ical Analysis — Molybdenum — General — Occurrence and Prep- 
aration — Molybdic Acid and the Molybdates — Lead Molyb- 
date — Phospho-moly bdic A cid — Tungsten — General — Occur- 
rence and Preparation — Properties — Tungstic Acid and the 
Tungstates — Uranium — General — Occurrence and Prepara- 
tion — Properties — Oxides — Uranous Salts — Uranyl Salts — 
Uranates. 

Experiments : Chromic Acid and the Chromates 593 



xvin CONTENTS, 



CHAPTER XXXII. 

ELEMENTS OF FAMILY VII, GROUP A : 

MANGANESE. 

PAGE 

General — Forms in which Manganese occurs in Nature — Prepar- 
ation and Properties — Manganous Chloride — General Re- 
marks concerning the Oxides— Manganous Hydroxide — 
Manganous-manganic Oxide — Manganic Oxide — Manganese 
Dioxide— Weldon's Process for the Regeneration of Man- 
ganese Dioxide in the Preparation of Chlorine — Sulphides — 
Manganous Sulphate — Manganic Sulphate — Manganic Acid 
and the Manganates — Permanganic Acid and the Perman- 
ganates — Potassium Permanganate — Reactions which are of 
Special Value in Chemical Analysis. 

Experiments : Manganese and its Compounds, 609 



CHAPTER XXXIII. 

ELEMENTS OF FAMILY VIII, SUB-GROUP A. 
IRON — COBALT — NICKEL. 

General — Iron — Introductory — Forms in which Iron occurs in 
Nature — Metallurgy — Varieties of Iron — Properties of Iron 
— Ferrous Chloride— Ferric Chloride — Cyanides — Potassium 
Ferrocyanide — Ferrohydrocyanic Acid — Ferric Ferrocyanide, 
or Prussian Blue — Potassium Ferricyanide — Ferrihydro- 
cyanic Acid— Ferrous Ferricyanide — Ferrous Hydroxide — 
Ferric Hydroxide — Ferrous-Ferric Oxide — Soluble Ferric 
Hydroxide — Ferric Oxide — Ferrous Sulphide — Ferrous Car- 
bonate — Ferrous Sulphate — Ferric Sulphate — Ferrous Phos- 
phate — Iron Disulphide — Reactions which are of Special 
Value in Chemical Analysis — Cobalt — General— Occurrence 
and Preparation — Properties — Cobaltous Chloride — Cobalt- 
ous Hydroxide — Cobaltic Hydoxide — Cobalt Sulphide — 
Cyanides — Smalt — Nickel — General — Occurrence and Prep- 
aration — Properties — Alloys — Other Applications of Nickel — 
Nickelous Hydroxide — Nickelic Hydroxide — Cyanides — 
Nickel Carbonyl — Reactions of Cobalt and Nickel which are 
of Special Value in Chemical Analysis. 

Experiments : Iron and its Compounds, 620 



CONTENTS. xix 



CHAPTER XXXIV. 



ELEMENTS OP FAMILY VIII, SUB-GROUP B : 

RUTHENIUM— RHODIUM— PALLADIUM. 

ELEMENTS OF FAMILY VIII, SUB-GROUP C : 

OSMIUM— IRIDIUM— PLATINUM. 



PAGE 



General-The Platinum Metals-Metallurgy-Ruthenium-Prep- 
aration -Properties - Osmium-Preparation-Properties- 
Rhodium-Iridium-Preparation-Properties-Palladium- 
Preparation-Properties-Palladium-Hydrogen-Platinum- 
Preparation-Properties-Applications of Platinum-Alloys 
of Platinum-Chlorides-Chlorplatinic Acid-Sulphides 

Experiments : Platinum— Chemical Analysis, ....'.. 641 

CHAPTER XXXV. 

SOME FAMILIAR COMPOUNDS OF CARBON. 
Organic Chemistry-Occurrence of the Compounds of Carbon- 
Formation of Hydrocarbons-Distillation of Coal— Distilla 
tion of Wood-Distillation of Bones-Fermentation-Classes 
of Compounds of Carbon - Compounds of Carbon and 
Hydrogen-Petroleum— Refining of Petroleum— Hydrocar- 
bons in Petroleum-Homology-The Ethylene Series of 
Hydrocarbons-The Acetylene Series-The Benzene Series 
-Marsh-gas, Methane, Fire-damp-Substitution-products of 
the Hydrocarbons-Chloroform— Iodoform— Ethylene Ole 
fiant Gas- Acetylene-Methyl Alcohol, Wood-spirit-Ethyl 
Alcohol, Spirits of Wine-What Change takes Place in 
Sugar ?-What Causes the Change ?-Germs in the Air- 
Different Kinds of Fermentation-Distillation of Fermented 
Liquids-Properties of Alcohol-Uses of Alcohol-Glycerin 
—Properties— Acetic Aldehyde, Ordinary Aldehyde-Chloral 
-Formic Acid-Acetic Acid-Properties-Uses-Salts of 
Acetic Acid-Fatty Acids-Butyric Acid-Palmitic Acid- 
Steanc Acid-Soaps-Use of Soap-Action of Soap on Hard 
Waters-Relations of the Soap Industry to Other Industries- 
Oxalic Acid-Lactic Acid-Malic Acid-Tartaric Acid-Citric 
Acid-Ether-Action of Acids upon Alcohols-Saponification 
-Fats-Butter-Ethereal Salts as Essences-Nitroglycerin 
Comparison of Formulas-Alcohols-More Complex Alcohols 
—Radicals or Residues— Acids. 
Experiments: Fermentation-Aldehyde-Soap-Hard Water,. 649 



xx CONTENTS. 

CHAPTER XXXVI. 

OTHER COMPOUNDS OF CARBON. 

PAGE 

Tlie Carbohydrates — Grape-sugar, Glucose, Dextrose — Formation 
of Dextrose — Manufacture of Dextrose or Glucose — Proper- 
ties — Levulose, Fruit-sugar — Cane-sugar — Sugar-refining — 
Molasses — Properties of Sugar — Sugar of Milk, Lactose — 
Souring of Milk — Cellulose — Properties — Gun-cotton, Py- 
roxylin, Nitrocellulose — Collodion — Celluloid — Paper — 
Starch — Manufacture of Starch — Properties — Flour — Bread- 
making — Aromatic Compounds — Nitrobenzene — Aniline — 
Aniline Dyes — Phenol, Carbolic Acid — Oil of Bitter Almonds, 
Benzoic Aldehyde — Benzoic Acid — Balsams and Odoriferous 
Resins— Gallic Acid — Tannic Acid, Tannin — Tanning — In- 
digo — Naphthalene — Anthracene — Alizarin — Glucosides — 
Myronic Acid — Alkaloids — Quinine — Cocaine — Nicotine — 
Morphine and Narcotine, 667 

Index 677 



A COLLEGE TEXT-BOOK OF 

CHEMISTRY, 



CHAPTER I. 

CHEMICAL AND PHYSICAL CHANGE. — EARLIEST 
CHEMICAL KNOWLEDGE.— LAW OF THE INDE- 
STRUCTIBILITY OF MATTER.— LAW OF DEFINITE 
PROPORTIONS.— LAW OF MULTIPLE PROPORTIONS. 
—THE ELEMENTS. 

Matter and Energy. — The sensible universe is made up 
of matter and energy. It is difficult to give satisfactory 
definitions of either of these terms, but, in a general way, 
it may be said that matter is anything which occupies 
space, and energy is that which causes change in matter. 
It requires but little observation to show that there are 
many kinds of matter, and apparently many kinds of 
energy. As examples of the different kinds of matter we 
have the many varieties of rocks and earth, as granite, 
limestone, quartz, clay, sand, etc. ; the plants and their 
fruits; the substances which enter into the composition of 
animals; and innumerable manufactured products. As 
examples of the different forms of energy, we have heat, 
light, motion, etc. Under the influence of the forms of 
energy the forms of matter are constantly undergoing 
change. Everywhere these changes are taking place. 
Changes in position and in temperature appeal most 
directly to our senses, and are most easily studied. But 
there are many other kinds of change which are of the 
highest importance. Thus there are electrical changes, 
manifestations of which we see in thunder-storms; there 



2 COLLEGE CHEMISTRY. 

are magnetic changes which may be studied to some extent 
by means of the magnetic needle; and there are, further, 
what are called chemical changes which affect the compo- 
sition of substances. 

Chemical Change. — For the purpose of study it is con- 
venient to distinguish between two classes of changes in 
matter, the difference between which can best be made 
clear by means of examples. Consider the changes in- 
cluded under the head of fire. We see substances destroyed 
by fire, as we say. They disappear as far as we can deter- 
mine by ordinary observation. When iron is exposed to 
the air a serious change takes place. It becomes covered 
with a reddish-brown substance which we call rust. If 
the piece of iron is comparatively thin, and it is allowed 
to lie in the air long enough, it is completely changed to 
the reddish-brown substance, and no iron as such is left. 
If the juices from fruits, as from apples, are allowed to 
stand in the air, they undergo change, becoming sour, and 
a somewhat similar change takes place in milk. If a spark 
is brought in contact with gunpowder there is a flash and 
the powder disappears, a dense cloud appearing in its 
place. 

In the changes referred to the substances changed dis- 
appear as such. After the fire, the wood or the coal, or 
whatever may be burned, is no longer to be found. The 
rusted iron is no longer iron. The gunpowder after the 
flash is no longer gunpowder. Changes of this kind in 
which the -substances disappear and something else is 
formed in their place are known as chemical changes. 

Physical Change. — There are many changes taking 
place that do not affect the composition of substances. 
Iron, for example, may be changed in many ways and still 
remain iron. It may become hotter or colder. Its posi- 
tion may be changed, or, as we say, it may be moved. 
The iron may be struck in such a way as to give forth 
sound. It may be made so hot that it gives light. When, 
for example, it becomes red-hot, it can be seen in a dark 
room. A piece of iron may be changed further by con- 



■■ 



PHYSICS AND CHEMISTRY. 3 

necting it with what is known as a galvanic battery. A 
current of electricity then passes through it, and we can 
easily recognize the difference between a piece of iron 
through which a current of electricity is passing and one 
through which no current is passing. The former when 
brought into certain liquids will at once change their com- 
position, while the latter will not cause such change. 
Finally, when a piece of iron is brought in contact with 
loadstone, it acquires new properties. It now has the 
power to attract and hold to itself other pieces of iron. In 
all these cases, the iron is changed, but it remains iron. 
After the moving iron comes to rest, it is exactly the same 
thing that it was before it was moved. After the iron that 
is giving forth sound has ceased to give forth sound, it 
returns to its original condition. Let the heated iron 
alone, and it cools down, soon ceasing to give light, and 
presenting no evidence of being warm. Eemove the iron 
from contact with the galvanic battery, and it loses those 
properties which are due to the current of electricity. In 
time, the iron which is magnetized by contact with the 
loadstone loses its magnetic properties. It then no longer 
has the power to attract other pieces of iron, and does not 
differ from ordinary iron. 

While iron has been taken as an example, other sub- 
stances undergo similar changes. These changes which 
do not affect the composition of the substances are called 
physical changes. 

Physics and Chemistry. — According to what has been 
said, we have two classes of changes presented to us for 
study : 

(1) Those which do not affect the composition of sub- 
stances, or physical changes. 

(2) Those which, do affect the composition of substances, 
or chemical changes. 

That branch of science which has to deal with physical 
changes is known as Physics. And that which has to 
deal with chemical changes is known as Chemistry. 

Everything that has to do with motion, heat, light, 



4 COLLEGE CHEMISTRY. 

sound, electricity, and magnetism is studied under the 
head of Physics. Everything that has to do with the 
composition of substances is studied under the head of 
Chemistry. It is, however, impossible to study these two 
subjects entirely independently of each other. Whenever 
a chemical change takes place, it is accompanied by 
physical changes; and in order that the former may be 
clearly understood, a study of the latter is necessary. 

Earliest Chemical Knowledge. — Those substances which 
are most abundant and most widely distributed in nature 
were, of course, the first known and studied ; and the same 
is true of those chemical changes which occur most com- 
monly and produce the most striking effects. Simply by 
observing those things which surround us and those 
changes in composition which take place naturally, a con- 
siderable amount of chemical knowledge might be gained, 
and indeed the earliest knowledge of chemistry was 
acquired in this way. It was not, however, until men came 
to experiment upon the substances which they found in 
nature, that knowledge of chemical changes made rapid 
progress. Since then an enormous amount of knowledge 
has been gained, and every year the stock is increased by 
new discoveries, until the field appears almost boundless. 

Law of the Indestructibility of Matter, — One of the 
first facts of fundamental importance established by the 
study of chemical phenomena is that after a chemical 
change has taken place there is just as much matter as 
there was before the change. Whenever matter apparently 
disappears, it continues to exist in some other form. If 
it were possible to annihilate matter or to call it into being, 
it would be of little or no value to weigh things. The 
first fundamental law bearing upon the changes in com- 
position which the different forms of matter undergo is 
the laiv of the indestructibility of matter. While, if we 
think of it, we can not conceive that this great law should 
not be true, we must not forget that the only way in which 
its truth could be established was by experiment. The 
law may be stated thus : 



CONSERVATION OF ENERGY. 5 

Whenever a change in the composition of substances takes 

place the amount of matter after the change is the same as 
before the change. 

According to this, and assuming that the law has always 
held good, it follows that the amount of matter in the uni- 
verse is the same to-day as it has been from the beginning. 
Transformations are constantly taking place, but these 
involve no increase nor decrease in the total amount of 
matter. 

Conservation of Energy. — Just as matter is neither 
created nor destroyed, so it has been made probable that 
the total amount of energy is unchangeable. One of the 
greatest discoveries in science was the recognition of the 
fact that one form of energy can be transformed into 
others, and that in these transformations nothing is lost. 
We now know that for a certain amount of heat we can get 
a certain amount of motion, and that for a certain amount 
of motion we can get a certain amount of heat. We know 
that a similar definite relation exists between heat and 
electrical energy, and between these and chemical energy. 
We know, for example, that a definite amount of heat can 
be obtained by burning a definite amount of a given sub- 
stance, and we know also that with a definite amount of 
heat we can produce a definite amount of chemical change. 
Modern investigation has shown that all the different 
forms of energy are convertible one into the other without 
loss. This important fact is generally spoken of as the 
law of the conservation of energy. Transformations of 
energy are taking place constantly, as transformations of 
matter are, but the total amount in each case remains the 
same. 

Composition of Matter. — The fact that first impresses 
one in studying the various forms of matter found in the 
earth is their great variety. We find an almost infinite 
number of kinds of matter, and the question at once sug- 
gests itself, of what are these things composed ? This 
question has long been asked, and it will be long before 
an entirely satisfactory answer is reached. Still, much 



6 COLLEGE CHEMISTRY. 

more is now known in regard to the subject than was known 
in past ages, and some progress is constantly being made 
towards a solution of the problem. 

Elements. — As chemical experimenting upon things 
advanced, the fact impressed itself more and more strongly 
upon investigators, that, of the large number of substances 
known, some can be converted into simpler ones by chem- 
ical action and some cannot. In other words, some sub- 
stances like water can be broken down by various methods 
into two or more others of different properties, and these 
when brought together again under proper conditions form 
the original substance. Water can be decomposed into 
two gases, hydrogen and oxygen. Elaborate experiments 
have shown that the weight of water decomposed is exactly 
equal to the weight of the hydrogen plus that of the 
oxygen obtained, and that when the hydrogen and oxygen 
are brought together again under proper conditions exactly 
as much water is formed as was originally decomposed. 
It appears, therefore, that water consists of at least two 
simpler substances. A similar conclusion is reached by a 
study of by far the largest number of the substances with 
which we have to deal. On the other hand, no treatment 
to which hydrogen and oxygen have been subjected has, 
as yet, effected their decomposition. They can be made 
to combine with other substances, as, for example, with 
each other, and thus form more complex substances, but 
nothing simpler than hydrogen has ever been obtained 
from hydrogen, and nothing simpler than oxygen has ever 
been obtained from oxygen. Whether the decomposition 
of these substances will ever be effected is a question that 
cannot be answered. All 'that we know is that at present 
they cannot "be decomposed. We therefore speak of them 
as elements, meaning by the term, that, with the means 
now at the disposal of chemists, it is impossible to get 
simpler substances from them. There are at present 
seventy-six substances known which are called elements 
for the same reasons that hydrogen and oxygen are called 
elements. It is quite possible that the number may be 



CHEMICAL ACTION. 7 

increased in the future, and it is also quite possible that 
the number may be decreased. New elements will in all 

probability be discovered, and probably some of the sub- 
stances now included in the list of elements will eventually 
be shown to be capable of decomposition. 

The view at present held in regard to the forms of 
matter which go to make up that part of the universe 
which comes under our observation is that they are all 
composed of the elementary substances. Many of them, 
like water, are composed of only two elements; others 
of three; and still others of four, five, six, and more; 
but most of them are comparatively simj)le, and rarely 
does any one contain more than four or five elements. 
Of the seventy elements known, only about twelve enter 
into the composition of most things with which we com- 
monly have to deal. The others occur in relatively small 
quantity. 

Chemical Action. — In the last paragraph it was stated 
that most substances can be decomposed, and that under 
proper conditions the elements combine. We must now 
inquire more carefully into the meaning of these expres- 
sions. Among the elements are the well-known substances 
lead, iron, and sulphur. If some finely-divided iron is 
brought in contact with sulphur, apparently no action 
takes place. If the two are put in a mortar and mixed, no 
matter how thoroughly, there is no evidence of action. 
The mixture has, to be sure, a different appearance from 
that of either constituent, but still both substances are 
present, and can be recognized in various ways. If, for 
example, a little of the mixture is examined with the aid 
of the microscope, particles of iron and of sulphur will be 
recognized lying side by side. If, further, the mixture is 
treated with the liquid, carbon disulphide, which has the 
power to dissolve the sulphur but not the iron, the sulphur 
will be dissolved while the iron will be left unchanged. 
Finally, if a dry magnet is introduced into the mixture, 
the iron will adhere to it, and by careful manipulation the 
two constituents can be separated. These facts furnish 



8 COLLEGE CHEMISTRY. 

evidence that both iron and sulphur are present in the 
mixture in unchanged condition, just as sugar and sand 
are present in a mixture of these two substances. If now 
the mixture of sulphur and iron is heated in a dry test- 
tube, marked changes will take place, and there will be 
formed a black substance entirely different from either of 
the elements employed in the experiment. Carbon disul- 
phide can no longer extract sulphur from it. The magnet 
can no longer pick out the iron, and under the microscope 
one homogeneous substance is seen instead of the two ele- 
mentary substances. If the experiment is performed with 
proper precautions, the amount of matter after the actum 
will be found to be exactly the same as before the action. 
A serious change has taken place, but no change in the 
amount of matter. The act is one of chemical combina- 
tion, and the substance formed is called a chemical com- 
pound. A few other examples will aid in making the 
conception of chemical combination clear. When a bit 
of phosphorus is brought in contact with a little iodine, 
action takes place at .once; the two elements combine, 
losing their own characteristic properties and forming a 
compound with properties quite different from those of 
the constituents. When the gases hydrogen and oxygen 
are brought together and a spark is passed through the 
mixture, an explosion occurs, and, in place of the gases, 
the liquid, water, is formed. When sulphur burns in the 
air the product formed is a pungent gas. It has been 
shown that the act consists in the combination of the sul- 
phur with the gas, oxygen, which is contained in the air. 
All these cases are examples of chemical combination. 
But chemical action may be of the opposite kind, that is 
to say, instead of being combination, it may be decomposi- 
tion. Thus, water which is formed by the chemical com- 
bination of hydrogen and oxygen may, by proper methods, 
be decomposed into the same elements. We may con- 
veniently think of that which causes elements to combine 
as an attractive force exerted between the elements. Now, 
when some power which can overcome this attraction is 



II. I 



CHEMICAL AFFINITY. 

brought to hear upon a compound, decomposition takes 
place, and the elements are, as we say. sei free. When, 
for example, the substance known as red oxide of mercury 

or mercuric oxide is heated to a sufficiently high tempera- 
ture a colorless gas is given off from it. and globules of 
mercury are formed at the same time. The gas, as will 
be shown later, is oxygen, so that from the red oxide of 
mercury, which is a chemical compound of mercury and 
oxygen, we get, by heating, the two elements in the free 
state. In this case, heat overcomes the chemical attrac- 
tion which, in the compound, holds the elements together. 
Chemical Affinity. — It is evident from what has already 
been said that there is some power which can hold ele- 
ments together in a very intimate way, so intimate that 
we cannot recognize them by ordinary means. We do not 
know what causes the sulphur and iron to combine, but 
w r e do know that they combine. Similarly, we do not 
know what causes a stone thrown in the air to fall back 
again, but we know that it falls back. It is true we may 
say that the cause of the falling of the stone is the attrac- 
tion of gravitation, but this does not give us any real 
information, for, if we ask what the attraction of gravita- 
tion is, we can only answer that it is that which causes all 
bodies to attract one another. AVe may also say, and do 
say, that the cause of chemical combination is chemical 
affinity. But in so doing we only give a name to something 
about which we know nothing except the effects it pro- 
duces. All the chemical changes that are taking place 
around us may, then, be referred to the operation of 
chemical affinity. If this power should cease to operate, 
what would be the result ? Nature would be infinitely less 
complex than it now is. All complex substances would 
be resolved into the elements of which they are composed, 
and, as far as we know, there would be only about seventy 
different kinds of substances. All living things would 
cease to exist, and in their place there would be three in- 
visible gases, and something very much like charcoal. 
Mountains would crumble to pieces, and all water would 



io COLLEGE CHEMISTRY. 

disappear, giving two invisible gases. The processes of life 
in its many forms would be impossible, as, however subtle 
that which we call life may be, we cannot imagine it to 
exist without the existence of certain complex forms of 
matter; and, as regards the life process of animals and 
plants, complex chemical changes are constantly taking 
place, within them, and these changes are essential to the 
continuance of life. 

Chemical Compounds and Mechanical Mixtures. — The 
substances formed by chemical combination of the ele- 
ments are called chemical compounds. Most substances 
found in nature are made up of several others. Wood, 
for example, is very complex, containing a large number 
of distinct chemical compounds intimately mixed together. 
Some of these can be isolated, but it is impossible to isolate 
them all with the means at present at our command. 
Most of the rocks met with are also quite complex, and it 
is difficult to isolate the constituents. If we look at a piece 
of coarse-grained granite, we see plainly enough that it 
contains different things mixed together, and if it is 
broken up we can pick out pieces of different substances 
from the mass. If we now examine a piece of each of the 
different substances thus picked out of the granite, it 
appears to be homogeneous, that is to say, we cannot 
recognize the presence of more than one kind of thing in 
any one piece. If the piece is carefully selected it may be 
powdered finely in an agate mortar, and some of the powder 
examined with a microscope without the presence of more 
than one substance being recognized. It is possible to 
isolate three substances from granite by simply breaking 
it up and picking out the pieces of different kinds. It 
might therefore be concluded that granite consists of three 
substances. This is true, but it is not the whole truth. 
For it is possible by proper means to get simpler sub- 
stances from each of the three already separated. This is, 
however, a much more difficult process than the separation 
first accomplished. To effect the separation of each of the 
three constituents of granite into its elements requires 



CHEMICAL COMPOUNDS AND CHANGES. II 

more powerful means. Substances must be brought in 
contact with them which act upon them, changing their 
composition, that is to say, act chemically upon them, and 
high heat must be used to aid the action. By skilful work 
it is, however, possible to separate the three components 
of granite into their elements. 

From the above it is evident that substances may be 
united in different ways. They may be so united that it 
is a simple thing to separate them by mechanical processes. 
Or they may be so united that it is impossible to separate 
them by mechanical processes. By a mechanical process 
is meant any process which does not involve the use of 
heat, electricity, or chemical change. Thus, the mechan- 
ical process made use of in the case of granite consisted in 
picking out the pieces. The separation of the particles 
of different sizes by means of a sieve is a mechanical 
process. The separation of two liquids which do not mix 
with each other is a mechanical process. Complex sub- 
stances which can be separated into their components by 
purely mechanical processes are called mechanical mix- 
tures. Thus granite is a mechanical mixture of three, 
chemical compounds. Similarly, most natural substances 
are more or less complex mixtures of chemical compounds, 
or, much more rarely, of elements. Air, ;for example, is 
a mechanical mixture consisting mainly of the two elements 
nitrogen and oxygen. It is not always an easy matter to 
distinguish between mechanical mixtures and chemical 
compounds, as there are mixtures which it is extremely 
difficult to resolve into their components, and there are, 
on the other hand, chemical compounds which are ex- 
tremely unstable. Generally, however, the difference is 
recognized without serious difficulty. 

Qualitative and Quantitative Study of Chemical 
Changes. — In general there are two ways in which chem- 
ical changes may be studied. Substances may be brought 
together under a variety of conditions and, if action takes 
place, the properties of the product or products may then 
be studied and compared with those of the substances 



12 COLLEGE CHEMISTRY. 

brought together. In the early periods of the history of 
chemistry the study was almost wholly of this kind. This 
is called qualitative study. But we may go farther than 
this, and take into consideration the weights or masses ol 
the substances we are dealing with. We should then be 
studying the changes quantitatively. By means of the 
quantitative method the law of the indestructibility of 
matter was placed upon a firm basis. By further use of 
this method other laws of the highest importance to the 
science of chemistry were soon brought to light. 

Law of Definite Proportions. — The fact that sulphur 
and iron combine chemically when a mixture of the two is 
heated has been referred to. The question whether they 
combine in all proportions is one that can be answered 
only by a quantitative study of the process. If the process 
were to be studied for the first time the method of pro- 
cedure would be this: We should mix the elements in 
different proportions and, after the action, we should 
determine whether any of either of the elements is left in 
the uncombined state; and, further, by decomposing the 
product, we should determine whether it always contains 
the elements in the same proportions. The problem, in 
this case, is by no means a simple one, but it has been 
repeatedly worked over with the greatest possible care, 
and, as the result of the work, the conclusion is justified 
that the product always contains the elements in exactly 
the same proportions. Similar work has been done for 
most other chemical compounds known, and the general 
conclusion known as the law of definite proportions has 
been drawn. This law may be stated thus: 

A chemical compound always contains the same constit- 
uents in the same proportion by weight. 

The truth of this general statement or law has not 
always been acknowledged by chemists. At the beginning 
of this century a celebrated discussion on the subject took 
place between. Proust and Berthollet. The discussion led 
to a great deal of careful work which tended to confirm the 
law, and since that time it' has not been seriously doubted. 



LA IV OF MULTIPLE PROPORTIONS. 1 3 

About twenty years ago a Belgian chemist, Stas, by a long 
series ot probably the most painstaking and accurate 
chemical experiments ever performed, showed that in the 
compounds which he worked with there was no variation 
in composition that could be detected by the most refined 
methods of chemistry. In the present state of our knowl- 
edge it appears that the law of definite proportions is a law 
in the strictest sense. 

Law of Multiple Proportions. — It does not require a 
verv extended study of chemical phenomena to show that 
from the same elements it is possible in many cases to get 
more than one product. Thus iron and sulphur form 
three distinct compounds with each other. Tin combines 
with oxygen in two proportions. The elements potassium, 
chlorine, and oxygen combine in four different ways, 
forming four distinct products. Nitrogen and oxygen 
form five products. In the early part of this century the 
English chemist, Dalton, by a study of cases like those 
mentioned was led to the discovery of another great law of 
chemistry known as the law of multiple proportions. 
Many substances had been analyzed before his time, and 
the percentages of the constituents determined with a fair 
degree of accuracy. He examined, first, two gases, both 
of which consist of carbon and hydrogen. He determined 
the percentages of the constituents, and found them to be 
as follows : 

Olefiant gas, 85.7 per cent carbon and 14.3 per cent 
hyd rogen. 

Marsh-gas, 75.0 per cent carbon and 25.0 per cent 
hydrogen. 

On comparing these numbers, he found that the ratio 
of carbon to hydrogen in olefiant gas is as 6 to 1 ; whereas 
in marsh-gas it is as 3 to 1 or 6 to 2. The mass of 
hydrogen, combined with a given mass of carbon, is 
twice as great in the one case as in the other. 

There are, further, two compounds of carbon and 
oxygen, and in analyzing these the following figures were 
obtained: 



14 COLLEGE CHEMISTRY. 

Carbon monoxide, 42.86 per cent carbon and 57.14 per 
cent oxygen. 

Carbon dioxide, 27.27 per cent carbon and 72.73 per 
cent oxygen. 

But 42.86 : 57.14:: 6 :8 and 27.27 : 72.73 ::6 : 16. 

The mass of oxygen combined with a given mass of car- 
bon in carbon dioxide is twice as great as the mass of oxy- 
gen combined with the same mass of carbon in carbon 
monoxide. These facts and other similar ones led to the 
discovery of the law of multiple proportions, which may 
be stated thus : 

If two elements A and B form several compounds with 
each other, and we consider any fixed mass of A, then the 
different masses of B which combine with the fixed mass of 
A bear a simple ratio to one another. 

By way of further illustration we may take the three 
compounds which iron forms with sulphur. In one of 
these, approximately 7 parts by weight of iron are in com- 
bination with 4 parts of sulphur; in a second, 7 parts of 
iron are in combination with 6 parts of sulphur; and in 
the third, 7 of iron are in combination with 8 of sulphur. 
The figures 4, 6, and 8 bear a simple ratio to one another 
which is 2 : 3 : 4. The five compounds of nitrogen and 
oxygen contain 7 parts by weight of nitrogen combined 
with 8, 16, 24, 32, and 40 parts of oxygen respectively. 
The figures representing the parts by weight of oxygen 
combined with 7 parts by weight of nitrogen are in the 
ratio 1:2:3:4:5. In the compounds formed by the 
elements chlorine, potassium, and oxygen the proportions 
by weight are represented in the following table : 



Chlorine. 


Potassium. 


Oxvpen. 


35.45 


39.15 


16 


35.45 


39.15 


32 


35.45 


39.15 


48 


35.45 


39.15 


64 



It will be observed that the ratio between the chlorine 
and potassium remains constant, but that the mass of 



COMBINING WEIGHTS OF THE ELEMENTS. 



•5 



oxygen varies regularly from 1(5 to 64; the masses bearing 
to one another the simple ratio 1:2:3:4. 

The law of multiple proportions like the law of definite 
proportions is simply a statement in accordance with what 
has been found true by experiment. Although discovered 
by Dalton at the beginning of this century and put forward 
upon what appears now to be only a slight basis of facts, 
all work since that time has confirmed it, and it forms 
to-day one of the corner-stones of the science of chemistry. 

Combining Weights of the Elements. — A careful study 
of the figures representing the composition of chemical 
compounds reveals a remarkable fact regarding the relative 
quantities of one and the same element which enter into 
combination with different elements. The proportions by 
weight in which some of the elements combine chemically 
with one another are stated in the following table : 



1 part 


H] 


•drogen combines 


with 


35.45 p 


arts Chlorine. 


1 " 




IC < ( 


" 


79.96 


" Bromine. 


1 " 




" 


" 


126.85 


" Iodine. 


35.45 part 


5 Chlorine combine with 


39.15 


" Potassium. 


79.96 


" 


Bromine " 


" 


39.15 


a u 


126.85 


<( 


Iodine " - 


" 


39.15 


a 


16 


i< 


Oxygen " 


< t 


65.4 


" Zinc. 


16 


" 


" " 


< < 


24.36 


" Magnesium 


16 


" 


it a 


" 


40 


" Calcium. 


16 


" 


" 


" 


137.4 


" Barium. 


65.4 


" 


Zinc 


< < 


32.06 


" Sulphur. 


24. 86 


" 


Magnesium " 


" 


32.06 


1! << 


40 


" 


Calcium " 


" 


32.06 


<< it 


137.4 


" 


Barium " 


< < 


32.06 


" 



It will be seen that the figures that express the relative 
weights of chlorine, bromine, and iodine that combine 
with 1 part by weight of hydrogen also express the relative 
weights of these elements that combine with 39.15 parts 
by weight ot potassium. So also the figures which express 
the relative weights of zinc, magnesium, calcium, and 
barium that combine with 10 parts by weight of oxygen 
also express the relative weights of these elements that 
combine with 32.06 parts by weight of sulphur. Now, an 



1 6 COLLEGE CHEMISTRY. 

examination of all compounds known has shown that 
hydrogen enters into combination with the other elements 
in the smallest proportions; it is therefore taken as unity 
in stating the relative weights of the other elements which 
enter into combination. The weight of another element 
that combines with 1* part by weight of hydrogen may be 
called its combining iveight. Thus, according to the above, 
the combining weights of chlorine, bromine, and iodine are 
respectively 35.45, 79.96, and 126.85. Similarly 39.15 is 
the combining weight of potassium, as it expresses the 
weight of potassium that combines with the above weights 
of chlorine, bromine, and iodine. Thus for every element 
a number can be selected, such that the proportions by 
weight in which the element enters into combination with 
others can be conveniently expressed by this number or 
by a simple multiple of it. These numbers are the com- 
bining weights. 

It is not by any means an easy matter to determine 
which numbers are most convenient for all circumstances; 
and if the selection is to be determined solely by conve- 
nience, there may be differences of opinion as to what is 
most convenient. We shall see a little later that, while 
the numbers primarily express the combining weights and 
nothing else, and are based solely upon a study of the 
composition of chemical compounds, they have come to 
have a deeper significance which will be explained in due 
time. Those adopted are called atomic weights because 
for strong reasons they are believed to express the relative 
weights or masses of the minute indivisible particles or 
atoms of which the various kinds of matter are assumed 
to be made up. The atomic theory, as it is called, will be 
treated of farther on, and the relation between the theory 
and the figures called the atomic weights will be discussed 
at some length. For the present it will be best to use the 
figures as expressing the combining weights, and as being 
entirely independent of any speculations regarding the 
constitution of matter and the existence of atoms. 

*Seep. 19. 



SYMBOLS AND ATOMIC WEIGHTS OF THE ELEMENTS. 1/ 

The Elements, their Symbols and Atomic Weights. — 

It has already been stated that there are about seventy 
elementary substances known, but that of these only a 
small number enter into the composition of common 
things to any great extent. It has been calculated that the 
solid crust of the earth is made up approximately as repre- 
sented in the subjoined table : 



Oxvgen 47.292 

Silicon 27.2135 

Aluminium 7.81$ 

Iron 5.46$ 



Calcium 



Sodium 2.36c; 

Potassium 2.40$ 



While oxygen forms a large proportion of the solid crust 
of the earth, it forms a still larger proportion (eight- 
ninths) of water, and about one-fifth of the air. Carbon 
is the principal element that enters into the structure of 
living things, while hydrogen, oxygen, and nitrogen also 
are essential constituents of animals and plants. Nitrogen 
forms about four-fifths of the air. 

In representing the results of chemical action, it is con- 
venient to use abbreviations for the names of the elements 
and compounds. Thus, instead of oxygen we may write 
simply 0; for hydrogen, H; for nitrogen, N; etc. These 
symbols are used in expressing what takes place when sub- 
stances act upon one another. Very frequently the first 
letter of the name is nsed as the symbol. If the names 
of two or more elements begin with the same letter, this 
letter is used, and some other letter of the name is added. 
Thus, B is ,the symbol of boron, Ba of barium, Bi of 
bismuth ; C is the symbol of carbon, Ca of calcium, Cd of 
cadmium, Ce of cerium, CI of chlorine, Cr of chromium, 
Cs of caesium, Cu of copper. In some cases the symbol is 
derived from the Latin name of the element. Thus, the 
S} r mbol of iron is Fe, from the Latin ferritin ; of copper 
Cu, from cuprum ; of mercury Hg, from hydrargyrum ; 
etc. 

The names themselves are formed in a variety of ways. 
Chlorine is derived from the Greek jAr*?po>, which means 
yellowish-green, as this is the color of chlorine. Bromine 



1 8 COLLEGE CHEMISTRY. 

comes from /3pc0jjos, a stench, a prominent characteristic 
of bromine being its bad odor. Hydrogen comes from 
vdoop, water, and yeveir, to produce, signifying that it 
is a constituent of water. Similarly nitrogen comes from 
virpov, niter, and yeveiv, to produce, nitrogen being 
one of the constituents of niter. Potassium is an element 
found in potash, and sodium is found in soda. Some 
elements have been named after the country in which they 
were first discovered. Thus we have gallium, discovered 
in France; scandium, discovered in Sweden; germanium, 
discovered in Germany. Tantalum was so called on 
account of the long-continued difficulties experienced in 
isolating it. Columbium received its name from the fact 
that it occurs in the mineral columbite, and this owes its 
name to the fact that it was first found in the United 
States of America. 

On page 19 is given a table containing the names of all 
the elementary substances now known, together with their 
symbols and atomic weights. The names of those which 
are most widely distributed, and form by far the largest 
part of the earth, are printed in small capitals. The 
names of those which are rare are printed in italics. 

The symbols of the elements represent not only the 
names but relative quantities. Thus stands for 10 parts 
by weight of oxygen; N for 14.01 parts by weight of 
nitrogen. Hydrogen enters into combination with other 
elements in the smallest relative quantity. It has the 
smallest combining weight, and it has been taken as the 
basis of the system. What the symbol really means then 
is that the weight of matter represented by it is 16 times 
as great as the weight of matter represented by the symbol 
H, and the weight of matter represented by the symbol N 
is 11.01 times as great as that represented by the symbol 
H; and so on through the list. Accurate investigation 
has shown, however, that the ratio between the combining 
weights of hydrogen and oxygen is not as 1 : 10, but as 
1:15.88. As many of the combining weights have been 
determined through the aid of the oxygen compounds of 



SYMBOLS AND ATOMIC WEIGHTS OF THE ELEMENTS. 19 



Element. Symbol. 

Aluminium Al 

Antimony Sb 

Argon A 

Arsenic As 

Barium Ba 

liisuiuth Bi 

Boron B 

Bromine Br 

Cadmium Cd 

CcBsium Cs 

Calcium Ca 

Carbon C 

Cerium Ce 

Chlorine CI 

Chromium Cr 

Cobalt Co 

Columbium Cb 

Copper Cu 

Erbium E 

Fluorine F 

Gallium Ga 

Germanium Ge 

Glucinum Gl 

Gold Au 

Helium \ . He 

Hydrogen H 

Indium In 

Iodine I 

Iridium Ir 

Iron Fe 

Krypton Kr 

Lanthanum La 

Lead Pb 

Lithium . . . Li 

Magnesium Mg 

Manganese Mn 

Mercury Hg 

Molybdenum Mo 



Atomic 
Weight, 

27.1 
120 

39.9 

75 
137.4 
208.5 

11 

79.96 
112.4 
133 

40 

12 
140 

35.45 

52.1 

59 

94 

63.6 
166 

19 

70 

72 
9.1 
197.2 
4 

1,01 
114 
126.85 
193 

56 
81.8 
138 
206.9 
7.03 

24.36 

55 
200.3 

96 



Element. Symbol 

Neodymium... , . . Nd 

Neon Ne 

Nickel Ni 

Nitrogen N 

Osmium Os 

Oxygen O 

Palladium Pd 

Phosphorus P 

Platinum Pt 

Potassium K 

Praseodymium.. . Pr 

Rhodium Rh 

Rubidium. ...... Rb 

Ruthenium Ru 

Samarium Sm 

Scandium Sc 

Selenium Se 

Silicon Si 

Silver Ag 

Sodium Na 

Strontium Sr 

Sulphur S 

Tantalum Ta 

Tellurium Te 

Thallium Tl 

Thorium Th 

Thulium Tu 

Tin ... , Sn 

Titanium Ti 

Tungsten W 

Uranium U 

Vanadium V 

Xenon X 

Ytterbium Yb 

Yttrium Y 

Zinc Zn 

Zirconium Zr 



Atomic 

M eight. 

143.6 
20 
58.7 
14.04 

191 
16 

106 
31 

194.8 
39.15 

140.5 

103 
85.4 

101.7 

150 
44.1 
79.1 
28.4 

107.93 
23.05 
87.6 
32.06 

183 

127 

204.1 

232.5 

171 

118.5 
48.1 

184 

239.5 
51.2 

128 

173 
89 
65.4 
90.7 



the elements, and calculated on the assumption that the 
combining weight of oxygen is 16, it is plain that, if we 
accept 15.88 as the combining weight of oxygen, many of 
the other combining weights must be changed accordingly. 
A large number of chemists, however, favor keeping the 
combining weight of oxygen 1G, and referring all the 
others to this as the basis of the system. This carries with 
it the awkward necessity of changing the combining weight 



20 COLLEGE CHEMISTRY. 

of hydrogen from 1 to 1. 01. For all ordinary purposes we 
may continue to use 1 as the combining weight of 
hydrogen. The above table is the one recently recom- 
mended by a German committee that has almost become 
international in its authority. 

There are very serious difficulties encountered in deter- 
mining the combining weights of the elements, and in 
regard to several given in the above table there is consider- 
able doubt as to their accuracy. Those of the elements 
with which we most frequently have to deal have, however, 
been determined with great care. Work in this field is 
being constantly carried on, and every year our knowledge 
in regard to the combining weights becomes more and 
more accurate. 

Symbols of Compounds. — As the elements enter into 
combination in the proportion of their respective combin- 
ing weights or simple multiples of these weights, it is an 
easy matter to represent the composition of compounds 
by means of the symbols. Thus hydrogen and chlorine 
combine in the proportion of their combining weights to 
form the compound hydrochloric acid. The compound is 
represented by the symbol HC1, which signifies that the 
compound contains hydrogen and chlorine in the propor- 
tion of 35.45 parts by weight of chlorine to 1.01 parts of 
hydrogen. So the symbol ZnO means a chemical com- 
pound consisting of 65.4 parts by weight of zinc and 16 
parts by weight of oxygen; HCN means a compound made 
up of 1.01 parts by weight of hydrogen, 12 parts by 
weight of carbon, and 14.04 parts by weight of nitrogen. 
Whenever the symbols of the elements are placed side by 
side with no sign between them, as in the above examples, 
the resulting symbol means that the elements are in 
chemical combination. But, as has been pointed out, 
elements may combine in more than one proportion. In 
one of the two compounds of carbon and oxygen the ele- 
ments are combined in the proportion of their combining 
weights, and the compound is represented by the symbol 
CO; in the other compound the elements are combined in 



SYMBOLS OF COMPOUNDS. 21 

the proportion of twice the combining weight of oxygen 
to the combining weight of carbon, and the compound is 
represented by the symbol 00 2 . The three compounds of 
iron and sulphur to which reference has already been made 
are represented by the symbols FeS, Fe 2 S 3 , and FeS 2 . 
The first represents a compound in which the elements are 
combined in the proportion of their combining weights, 
or 56 parts iron to 32. OG parts sulphur; the second repre- 
sents a compound in which the elements are combined in 
the proportion of twice the combining weight of iron 
(2 X 56 = 112 parts) to three times the combining weight 
of sulphur (3 x 32.04 = 96.12 parts); and the third 
represents a compound in which the elements are com- 
bined in the proportion of the combining weight of iron 
(56 parts) to twice the combining weight of sulphur (2 x 
32.04 = 64.08 parts). The four compounds of potassium, 
chlorine, and oxygen above mentioned are represented by 
the symbols KCIO, KC10 2 , KC10 3 , and KC10 4 , the 
meaning of which will be clear from the explanation just 
given. By means of such symbols all chemical compounds 
can be represented, and they show not only what ele- 
ments are contained in the compounds, but in what pro- 
portions the elements are combined. They represent 
facts that have been determined by experiment. Knowing 
the actual weight of one constituent of any compound we 
can calculate by the aid of the symbol the actual weights 
of the other constituents and of the compound itself. 
Thus, if we know the actual weight of the chlorine con- 
tained in a quantity of potassium chlorate, KC10 3 , we can 
calculate how much potassium and how much oxygen are 
contained in that same quantity, and also w r hat the quantity 
of potassium chlorate is. Suppose, for example, we know 
that in a certain quantity of potassium chlorate there is 
contained 25 grams of chlorine, and it is desired to know 
how much potassium and how much oxygen there are in 
this quantity, and also what the quantity of potassium 
chlorate is. We know that the compound KC10 3 is made 
up of 39.15 parts by weight of potassium, 35.45 parts by 



22 COLLEGE CHEMISTRY. 

weight of chlorine, and 3 times 16, or 48 parts by weight 
of oxygen, the whole making 122.6 parts by weight. The 
solution of the following equations in proportion will give 
the quantities desired: 

35.45 : 25 :: 39.15 : weight of potassium; 

35.45 :25 :: 48 : " " oxygen; 

35.45 : 25 :: 122.6 : " " potassium chlorate. 

Chemical Equations. — In dealing with cases of chemical 
action it is desirable to express by means of the symbols 
which represent the elements and compounds what takes 
place. In general, a chemical change is called a chemical 
reaction, and these reactions are of three kinds : 

(1) Two or more elements or compounds may unite 
directly to form one product. This is called combination. 
The following examples will suffice. When mercury is 
kept boiling in the air for a time it becomes covered with 
a layer of a red substance which is a compound of mercury 
and oxygen represented by the symbol HgO. Magnesium 
burns in the air and forms the compound MgO. Hydro- 
chloric acid, HC1, combines directly with ammonia, NH 3 , 
and forms the compound known as ammonium chloride, 
NH 4 CL Water, H 2 0, combines directly with lime or cal- 
cium oxide, CaO, to form slaked lime or calcium hy- 
droxide, Ca0 2 H 2 . To represent these facts, the symbols 
of the elements or compounds that act upon each other are 
written with a plus sign between them, and the sign of 
equality is written before the symbol of the product. The 
chemical equations that represent the above-mentioned 
chemical reactions are: 

Hg +0 =HgO; 

Mg + = MgO ; 
HC1 + NH 3 = NH 4 C1; 
H 2 + CaO = Ca0 2 H 2 . 

In reading the equations the plus sign is generally rendered 
by and, and the sign of equality by give. The first equa- 



«M^- 



CHEMICAL EQUATIONS. 2$ 

tion should accordingly be read, "Mercury and oxygen 
give mercuric oxide;" but it represents besides this fact 
the exact relations by weight between the quantities of the 
elements and of the compound that take part in the reac- 
tion. 

(2) The second kind of chemical reaction is decomposi- 
tion or the opposite of combination. Examples are fur- 
nished by the decomposition of mercuric oxide into mercury 
and oxygen by heat; of potassium chlorate into potassium 
chloride and oxygen by heat; of water into hydrogen and 
oxygen by the electric current; and of calcium carbonate 
or limestone, CaC0 3 , into lime or calcium oxide, CaO, 
and carbon dioxide, C0 2 , by heat. These reactions are 
represented by the following equations : 

HgO =Hg +0; 
KC10 3 = KC1 +30; 
H 2 =2H +0; 

CaC0 3 = CaO + C0 2 . 

The expressions 30 and 2H mean respectively three 
times the combining weight of oxygen and twice the 
combining weight of hydrogen, the figure being generally 
used in this way when the element is not in combina- 
tion. It is, however, sometimes w r ritten the same as if 
the element were in combination, as will be explained 
later. 

(3> The third kind of chemical reaction is that in which 
two or more compounds give rise to the formation of two 
or more others; or an element and a compound may act 
in such a way as to give another compound and another 
element. This is called double decomposition or metathesis. 
The following cases will serve as examples: Sulphuric 
acid, H 2 S0 4 , acts upon potassium nitrate, KNO H , or salt- 
peter, forming potassium sulphate, K 2 S0 4 , and nitric acid, 
HN0 3 ; nitric acid, HN0 3 , acts upon sodium carbonate, 
Na 2 C0 3 , forming sodium nitrate, NaN0 3 , carbon dioxide, 
C0 2 , and water, H 2 0; hydrochloric acid, HC1, and zinc, 



24 COLLEGE CHEMISTRY. 

Zn, give zinc chloride, ZnCl 2 , and hydrogen. These facts 
are 

H 9 S0, + 2KN0, = K 9 SO, + 2HN0. 

H 2 0; 



give zinc chloride, ZnCl 2 , and hydi 
represented as below : 

H 2 S0 4 + 2KN0 3 = K 2 S0 4 + 2HN0 3 ; 
2HN0 3 + Na 2 C0 3 = 2NaNO s + C0 2 + I 
2HC1 + Zn = ZnCl 2 + 2H. 



In the expressions 2KN0 3 , 2HN0 3 , 2NaN0 3 , and 
2HC1, the large^ figures placed before the symbol of the 
compounds signify that the quantities of the compounds 
represented by the symbol are to be multiplied by the 
figure. Thus, HC1 stands for 1.01 + 35.45 = 36.46 parts 
of hydrochloric acid; but 2HC1 stands for 2(1.01 -f 35.45) 
= 72.92 parts; 3HC1 stands for 3(1.01 + 35.45) = 109.38 
parts; etc. 

The Scope of Chemistry. — A complete study of chemis- 
try would involve the study of the action of all the ele- 
ments upon one another under all possible circumstances, 
and a study of the action of all compounds upon one 
another and upon the elements under all circumstances. 
This indicates that the field is almost boundless; and if 
the facts were not related among one another, if every 
time a reaction is studied we are to expect something 
entirely different from all other reactions, the task would 
be practically hopeless. Fortunately a great many gen- 
eral facts are known, and reactions which at first seem to 
have no connection are by careful study shown to be 
related. Thus the study is very much simplified and made 
interesting. It must be our purpose to study the facts in 
as systematic a way as possible, and to be constantly on 
the alert to detect relations. The habit of comparing a 
new reaction with others already studied should be culti- 
vated. In this way light will come out of the darkness, 
and the subject will gradually become clear. While the 
.simplest way to begin the study of chemistry is by a con- 
sideration of the elements, the subject is complicated by 
the fact that we cannot readily obtain these elements with- 
out the aid of substances that have not been studied, and 
of processes that are incomprehensible. There are, how- 



CHEMICAL REACTIONS. 25 

ever, two elements that occur in nature in enormous 
quantities, that can be obtained in the uncombined condi- 
tion quite easily. As the kinds of action which they 
exhibit are of great importance and are well calculated to 
give an insight into the nature of chemical action in 
general,, we may profitably begin our study of chemical 
phenomena by a study of these two elements. They are 
oxygen and hydrogen. In learning the main facts in 
regard to these two elements we shall learn a great deal 
that will be of importance in enabling us to understand 
other chemical phenomena; we shall learn how to study 
things chemically; and we shall thus prepare ourselves for 
a systematic study of the science of chemistry. 

Chemical Action accompanied by other kinds of Action. 
— Whenever a chemical change takes place it is accom- 
panied by other changes and, in order to gain a complete 
knowledge of the phenomena, these other changes must 
be studied. Thus when sulphuric acid acts upon zinc the 
chemical change is represented both qualitatively and 
quantitatively by the equation 

Zn -f H 2 S0 4 = ZnS0 4 + 2H. 

In studying the reaction, the first thing to do is to learn 
the nature of the substances formed, and the relations 
between the substances that act upon each other and the 
products. This may be called the purely chemical study 
of the reaction. But much more can be learned in regard 
to it by careful observation. In the first place, we must 
take into account the fact that a solid and liquid here react 
to form a solid and a gas; and the question suggests itself, 
does this change to the gaseous condition exert any influ- 
ence on the reaction, or is this a fact of no special impor- 
tance ? Again, it will be observed that accompanying the 
chemical change there is a marked rise in temperature, and 
we naturally inquire whether the quantity of heat evolved 
is definite for definite quantities of the substances, and, if 
so, what relation exists between them. There are still 



26 COLLEGE CHEMISTRY. 

other changes which must be taken into account in order 
to get a complete knowledge of a chemical reaction, but, 
as yet, the study of the other changes has not been taken 
up in a general way, and our information in regard to them 
is comparatively limited. Within late years much progress 
has been made in the study of the heat changes which accom- 
pany chemical changes. It has been found that every 
chemical change gives rise either to an evolution or to an 
absorption of heat, and that for definite quantities of the 
same substances under the same circumstances the same 
amount of heat is evolved or absorbed. The special study 
of the heat changes connected with chemical changes is 
called thermochemistry. A consideration of the facts and 
laws of thermochemistry is of assistance in dealing with 
chemical reactions, and some attention will be paid to the 
subject in this book. 

EXPERIMENTS. 

Chemical Change Caused by Heat. 

Experiment 1. — In a clean, dry test-tube put enough white 
sugar to make a layer i to i an inch thick. Hold the tube in 
the flame of a spirit-lamp or a laboratory burner. What evidence 
is furnished by this experiment that chemical change may be 
caused by heat ? What is left in the tube? Is it soluble? Is it 
sweet ? Is it sugar ? 

Experiment 2.— Half-fill the bulb of an arsenic-tube with red 
oxide of mercury, or, if such a tube is not available, proceed as 
follows : From a piece of hard-glass tubing of about 6 to 7 milli- 
metres (J inch) internal diameter cut off a piece about 10 centi- 
metres (4 inches) long by making a mark across it with a triangu- 
lar file, and then seizing it with both hands, one on each side of 
the mark, pulling and at the same time pressing slightly as if to 
break it. Clean and dry it, and hold one end in the flame of a 
blast-lamp until it melts together. During the melting turn the 
tube constantly around its long axis so that the heat .lay act 
uniformly upon it. Put into the tube thus made enough red oxide 
of mercury (mercuric oxide) to form a layer about 12 millimetres 
(-£ inch) thick. Heat the tube as in the last experiment. What 
change in color is noticed ? What is deposited upon the glass in 



CHANGES EFFECTED BY ELBCTRIC CURRENT. 



27 



the upper part of the tube ? What evidence is furnished by this 
experiment that chemical change can be effected by heat ? 



Chemical Changes can" be effected by an Electeic 
Cue rent. 

Experiment 3. — To the ends of insulated copper wires connected 
with two cells of a Bunsen's or Grove's battery fasten platinum 
plates, say 25 mm. (1 inch) long by 12 mm. (£ inch) wide. Insert 
these platinum electrodes into water contained in a shallow 
glass vessel about 15 cm. (G inches) wide and 7 to 8 cm. (3 inches) 
deep, taking care to keep them separated from each other. No 
action will take place, for the reason, as has been shown, that 
water will not conduct the current, and hence when the platinum 
electrodes are kept apart there is no current., By adding to the 
water about one tenth its own volume of strong sulphuric acid, it 
acquires the power to convey the eurrent. It will then be ob- 
served that bubbles rise from each of the platinum plates. In 
order to collect them an apparatus like 
that shown in Fig. 1 may be used. 

h and represent glass tubes which 
may conveniently be about 30 cm. (1 foot) 
long and 25 mm. (1 inch) internal diam- 
eter. They are first filled with the 
water containing one tenth its volume 
of sulphuric acid, and then placed with 
the mouth under water in the vessel A. 
The platinum electrodes are now brought 
beneath the inverted tubes. The bubbles 
which rise from them will pass upward in 
the tubes, and the water will be pressed 
down. Gradually the water will be com- 
pletely forced out of one of the tubes, 
while the other is still half full of water. 
The substances thus collected in the 
tubes are invisible gases. After the first 
tube is full of gas, place the thumb over its mouth and remove 
the tube. Turn it mouth upward and at once apply a lighted 
match to it. A flame will be noticed. The gas which was con- 
tained in the tube is therefore capable of burning. It cannot, 
therefore, have been air. In the meantime the second tube will 
have become filled with gas. Remove this tube in the same way 
and insert a thin piece of wood with a spark on it. The spark 




Fig. 1. 



28 COLLEGE CHEMISTRY. 

will at once burst into flame, and the burning of the wood will 
take place more actively than it does in ordinary air, as may be 
shown by withdrawing it and again inserting it into the tube. 
The gas in this tube, it will be noticed, does not take fire. With- 
out going into further details, it is clear from the above experi- 
ment that when an electric current acts on water two invisible 
gases are produced. A chemical change is caused by an electric 
current. 



Mechanical Mixtures and Chemical Compounds. 

Experiment 4. — Examine carefully a piece of coarse-grained 
granite ; break off some of it, and separate the constituents. 
How many are there ? By what properties do you recognize 
them ? Powder a small bit of one of the constituents, and exam- 
ine the powder with the microscope. Do you recognize more 
than one kind of matter ? Mix the powder of the three constitu- 
ents, and see whether in the mixed powder there is any difficulty 
in detecting the three kinds of matter with the aid of the micro- 
scope. 

Experiment 5. -Mix a gram or two of powdered roll-sulphur 
and an equal weight of very fine iron filings in a small mot tar. 
Examine a little of the mixture with a microscope. Not only can 
we recognize the particles of iron and of sulphur by means of the 
microscope, but we can also pick out the pieces of iron by means of 
a magnet. The magnet attracts the iron but not the sulphur, so 
that by passing the magnet often enough through the mixture we 
can pick out all the iron and leave all the sulphur. This separa- 
tion is really a mechanical separation. It is only a somewhat 
more refined method of picking out than that used in the case of 
granite. 

Experiment 6. — Pass a small magnet through the mixture above 
prepared. Unless 1he substances used are thoroughly dry, parti- 
cles of sulphur will adhere to the magnet, but even then it will be 
seen that most of that which is taken out of the mixture is iron. 

The iron and sulphur can also be separated by treating the mix- 
ture with a liquid known as carbon disulphide. Sulphur dissolves 
in this liquid, but iron does not. So that when the mixture is 
treated with it the iron is left behind, and can easily be recognized 
as such. 

Experiment 7. — Pour a few cubic centimetres of carbon disul- 
phide on a little powdered roll-sulphur in a dry test-tube. What 
takes place ? Treat iron filings in the same way. What takes place ? 



VARIOUS EXAMPLES OF CHEMICAL ACTION. 29 

Now treat a small quantity of the mixture with carbon disulphide. 
After the sulphur is dissolved filter the solution upon a good-sized 
watch-glass and let it stand. Examine what remains undissolved 
in the test-tube, and satisfy yourself that it is iron. After the 
liquid has evaporated examine what is left in the watch-glass and 
satisfy yourself that it is sulphur. Why are you justified in con- 
cluding that the substance left in the test-tube is iron and that 
left on the watch-glass is sulphur ? 

Experiment 8. — Make a fresh mixture of three grams each of 
powdered roll-sulphur and fine iron filings. Grind them together 
intimately in a dry mortar and put them in a dry test-tube. 
Heat gradually until the mass begins to glow. At first the 
sulphur melts and becomes dark-colored. It may even take fire. 
But soon something else evidently takes place. The whole mass 
begins 'to glow, and if you at once take the tube out of the flame, 
the mass continues to glow, becoming brighter. This soon stops ; 
the mass grows dark and gradually cools down. As soon as it 
reaches the ordinary temperature, the tube should be broken and 
the contents put in a mortar. A close examination will show that 
the mass does not look like the mixture of sulphur and iron 
with which we started. It has a bluish-black color, and is appar- 
ently homogeneous. An examination with the microscope, the 
magnet, and carbon disulphide will prove that, while there may 
be a little iron left, and possibly a little sulphur, most of the 
bluish-black mass is neither iron nor sulphur, but a new sub- 
stance with properties quite different from those of iron and from 
those of sulphur. 



Other Examples of Chemical Action. 

Experiment 9. — Examine a piece of calc-spar or marble. You 
see that it is made up of pieces of definite shape. It is, as we say, 
crystallized. It is quite hard, though a knife will cut it. Heated 
in a hard-glass tube, as in Experiment 2, it does not melt, but 
remains essentially unchanged. It does not dissolve in water. 
To prove this, put a piece the size of a pea in a test-tube with 
pure water. Thoroughly shake, and then, as heating usually aids 
solution, boil. Now pour off a few drops of the liquid on a piece of 
platinum-foil or an evaporating-dish, and by gently heating cause 
the water to evaporate. If there is anything in solution, there will 
be a solid residue on the platinum-foil or watch-glass. If not, 
there will be no residue. Now treat a small piece of the substance 



3° 



COLLEGE CHEMISTRY. 



with dilute hydrochloric acid and notice what takes place. Bub- 
bles of gas are given off. After the action has continued for 
about a minute, insert a lighted match in the upper part of the 
tube. It is extinguished, and the gas does not burn. The gas 
formed in this case is therefore plainly not identical with either 
one of those obtained from water by the action of the electric 
current (see Experiment 3). It is what is commonly called car- 
bonic-acid gas. As the action continues, the piece of calc-spar or 
marble grows smaller and smaller, and finally disappears, when 
there is a clear solution. The substance has dissolved in the hydro- 
chloric acid. In order to determine whether anything else has 
taken place besides the dissolving, we shall have to get rid of the 
excess of hydrochloric acid. This we can easily do by boiling it, 
when it passes off in the form of vapor, and then whatever is in 
solution will remain behind. For this purpose, put the solution 

in a small, clean porcelain evapo- 
rating-dish, and put this on a 
vessel containing boiling water, 
or a water-bath. The operation 
should be carried on in a place 
in which the draught is good, so 
that the vapors will not collect 
in the working-room. They are 
not poisonous, but they are an- 
noying. The arrangement for 
evaporating is represented in 
Fig. 2. 

After the liquid has evapo- 
Fig. 2. rated and the substance in the 

evaporating-dish is dry, examine it and carefully compare its 
properties with those of the substance which was put into the 
test-tube. Its structure will be found not to present the regulari- 
ties noticed in the original substance. It is much softer, dissolves 
in water, melts when heated in a hard-glass tube. It does not give 
off a gas when treated with hydrochloric acid. When exposed to 
the air it soon becomes moist, and after a time liquid. The experi- 
ment shows that when hydrochloric acid acts upon calc spar or 
marble the latter at least loses its own properties. It might be 
shown that some of the hydrochloric acid also loses its properties. 
In place of the two we get a new substance with entirely differ- 
ent properties. The two substances have acted chemically upon 
each other and produced a chemical compound. In this case it 
was only necessary to bring the substances in contact in order to 




VARIOUS EXAMPLES OF CHEMICAL ACTION. 3 1 

cause them to act chemically upon each other. It was not neces- 
sary to heat them, as it was in the case of the iron and sulphur. 

Experiment 10. — Bring together in a test-tube a few small 
pieces of copper and some moderately dilute nitric acid. In a 
short time action begins. The upper part of the tube becomes 
filled with a dark, reddish-brown gas which has a disagreeable 
smell. Do not inhale it, as when taken into the lungs it produces 
bad effects. The solution becomes colored dark blue, and the 
copper disappears. Examine this solution, as in Experiment 9, 
and see what has been formed. What are the properties of the 
substance found after evaporation of the liquid ? Is it colored ? 
Is it soluble in water ? Does it change when heated in a tube ? 
Is it hard or soft ? Does it in any way suggest the copper with 
which you started ? 

Experiment 11. — Try the action of dilute sulphuric acid on a 
little zinc in a test-tube. A gas will be given off. Apply a lighted 
match to it. Does the result suggest anything noticed in an ex- 
periment already performed ? After the zinc has disappeared, 
evaporate the solution as in Experiment 9. Carefully compare the 
properties of the substance left behind with those of zinc. 

Experiment 12. — Hold the end of a piece of magnesium ribbon 
about 20 centimetres (8 inches) long in a flame until it takes fire ; 
then hold the burning substance quietly over a piece of dark 
paper, so that the light, white product may be collected. Compare 
the properties of this white product with those of the magnesium. 
Here again a chemical act has taken place. The magnesium has 
combined with something which is found in the air, and heat was 
produced by the combination. The product is the white sub- 
stance. 

Experiment 13.— In a small, dry flask (400 to 500 ccm.) put a 
bit of granulated tin. Pour upon it 2 or 3 ccm. concentrated 
nitric acid. If no change takes place, heat gently, and presently 
there will be a copious evolution of a reddish-brown gas with a 
disagreeable smell (under what conditions has a gas like this 
already been obtained ?), the tin will disappear, and in its place 
will appear a white powder. Compare the properties of this white 
powder with those of tin. Why are you justified in concluding 
that they are not the same thing ? 



CHAPTEE II. 
A STUDY OF THE ELEMENT OXYGEN. 

Historical. — The older chemists considered air to be a 
simple substance, but the experiments of Priestley (1774) 
and of Scheele (1775) showed that the ai.r contains two 
gases, only one of which has the power to support combus - 
tion; and they succeeded independently of each other in 
showing that oxygen is a distinct substance. The discovery 
of oxygen had a very important bearing on the work of 
Lavoisier on combustion, and it was he who gave the name 
oxygen (or oxygene) to the gas, for the reason that he sup- 
posed it to be the essential constituent of all those chem- 
ical substances which are known as acids, the word being- 
derived from the Greek oSivcr, acid, and yeveiv, to pro- 
duce. While this is generally true, it has since been found 
that the acid properties of substances are not dependent 
upon oxygen, and, therefore, the name is misleading. 

Occurrence. — Oxygen is the most widely distributed and 
most abundant element of the earth. It forms, as has 
been stated, from 44 to 48 per cent of the solid crust of 
the earth; eight-ninths of water; and about one-fifth of 
the air. It occurs also in combination with carbon and 
hydrogen, or with carbon, hydrogen, and nitrogen in the 
substances which go to make up the structure of living 
things, whether vegetable or animal. Besides this it forms 
a part of most manufactured chemical products. 

Preparation. — Notwithstanding the abundant supply of 
oxygen in nature it is not a simple matter to get it in the 
free or imcombined state from most substances found in 

32 






OXYGEN— PREPARATION. 33 

nnt urc. As it forms eight-ninths of water, and water con- 
sists of only hydrogen and oxygen, the idea suggests itself 
at once that it may be made by the decomposition of 
water. This can be accomplished without serious difficulty 
by means of an electric current, and both hydrogen and 
oxygen can be obtained in this way; but the method is 
expensive and more complicated than others which are 
available. In the air the two gases nitrogen and oxygen 
are mixed together in the proportion of 1 volume of oxygen 
to 4 volumes of nitrogen. Here then, too, as in water, 
we have an enormous supply, but it is difficult to separate 
the oxygen from the nitrogen in such a way as to leave it 
uncombined. This can, however, be accomplished, and 
a method is now in practical use on the large scale for the 
purpose of preparing oxygen from the air. The method 
is based upon the fact that when barium oxide, BaO, is 
heated in a current of air it takes up oxygen and is con- 
verted into barium dioxide, Ba0 2 ; and when the pressure 
upon the dioxide is sufficiently diminished, it is decom- 
posed into the oxide, BaO, and oxygen, as represented in 
the equation 

Ba0 2 = BaO + 0. 

By this means oxygen can be extracted from the air and 
obtained in the free condition. 

Air can now be liquefied on the larger scale. When the 
liquid is allowed to stand in an open vessel at ordinary 
temperatures, the nitrogen boils off before the oxygen and 
a part of the latter is left behind. This method may be 
used for the preparation of oxygen from the air. 

Among natural substances that can be used for the 
preparation of oxygen, manganese dioxide or pyrolusite, 
also called the black oxide of manganese, Mn0 2 , is the 
most important. It gives off a part of its oxygen when 
heated to a comparatively high temperature. It has been 
shown that the decomposition is represented by this 
equation : 

3Mn0 2 = Mn 3 4 + 20. 



34 COLLEGE CHEMISTRY. 

As will be seen, only one-third of the oxygen contained 
in the dioxide is thus obtained in the free state. A similar 
method is that used by Priestley when he discovered 
oxygen. It consists in heating mercuric oxide, HgO, 
when it is decomposed as represented thus : 

HgO = Hg + 0. 

The substance potassium chlorate, KC10 3 , is manufac- 
tured on the large scale for a variety of purposes and is, 
therefore, easily obtained. It gives up its oxygen when 
heated. At first the decomposition represented by this 
equation takes place : 

SKCIO, = 5KC10 4 + 3KC1 + 40. 

The products are potassium pcrchlorate. KC10 4 , potas- 
sium chloride, KC1, and oxygen, one-sixth of the total 
oxygen being given off in the first stage. This part of 
the decomposition takes place readily, and at a compara- 
tively low temperature. If, after it is complete, the tem- 
perature is raised considerably higher, more gas is given 
off and the change represented by the equation 

KC10 4 = KOI + 40 

is accomplished. The final result is, therefore, the setting 
free of all the oxygen contained in the chlorate. This fact 
is represented thus : 

KOIO3 == KC1 + 30. 

The best method for use in the laboratory for the prep- 
aration of oxygen consists in heating a mixture of equal 
parts of coarsely-powdered manganese dioxide and potas- 
sium chlorate. This mixture gives up oxygen readily 
under the influence of heat. Potassium chlorate alone 
requires to be heated to a temperature of over 350° C. to 
effect its decomposition, but when mixed with manganese 
dioxide the decomposition takes place at about 200° C. 
The manganese dioxide does not lose any of its oxygen 
under the circumstances. Other substances, such as 



PHYSICAL AND CHEMICAL PROPERTIES. 35 

ferric oxide, copper oxide, etc., may be used with similar 
effect. 

Physical Properties. — Oxygen is a colorless, tasteless, 
inodorous gas. It is only slightly soluble in water, 100 
volumes of water at 0° dissolving 4. 1 volumes of oxygen. 
It is slightly heavier than air, its specific gravity is 
1.1056, and 1 litre under 760 mm. pressure and at tem- 
perature 0° weighs 1.429 grams, while a litre of air weighs 
1.2932 grams. In dealing with chemical elements and 
compounds which are gaseous, it is customary to use 
hydrogen instead of the air as the standard for specific 
gravity. While the specific gravity of oxygen in terms 
of air is 1.1056, in terms of hydrogen it is 15.89, or, in 
other words, a given volume of oxygen weighs 15.89 as 
much as the same volume of hydrogen under the same 
conditions. Under a pressure of 50 atmospheres and at a 
temperature — 119° it is condensed to a liquid of specific 
gravity 0.978. Liquid oxygen is a pale steel-blue trans- 
parent and very mobile liquid which boils at — 181° at 
ordinary pressure. When the pressure is reduced or 
removed, evaporation takes place so rapidly that a part of 
the oxygen is often frozen to a white solid. 

Chemical Properties. — At ordinary temperatures oxygen 
does not act readily upon most other things, as can be 
clearly shown by putting a variety of substances in the gas 
without heating them. If they are left for a considerable 
time some evidence of change may be observed, but gen- 
erally the change is extremely slow unless the temperature 
is raised. At higher temperatures, different for different 
substances, it combines with all the elements except 
fluorine, and it acts readily upon a large number of com- 
pounds. Its action is generally accompanied by an evolu- 
tion of heat and light, and the process under these 
circumstances is called combustion. This action may be 
illustrated by first heating and then introducing into 
vessels containing oxygen, sulphur, charcoal, iron in the 
shape of a steel watch-spring, and a bit of phosphorus. 
The phenomena observed show that chemical action takes 



3 6 COLLEGE CHEMISTRY. 

place, but they do not show what is formed. It is evident 
that in each case light and heat are evolved, and that the 
substances introduced into the oxygen are changed to 
other things. In the case of phosphorus the light given 
off is very intense, while in that of carbon and that of 
sulphur it is only slight. In the vessel in which the 
burning of the iron takes place a reddish-brown substance 
is deposited, while in that in which the phosphorus is 
burned dense white fumes are formed and at first the 
product is partly deposited upon the walls of the vessel in 
the form of a white powder that looks like snow. After 
standing for some time over water it disappears and the 
water evidently contains something in solution. A 
thorough study of the reactions above mentioned has 
shown that they consist in the chemical combination of 
oxygen with the substances burned. The light and heat 
are results of the chemical action. The reactions are 
represented by the following equations : 

With sulphur, S -f- .20 = S0 2 ; 

" carbon, + 20 = C0 2 ; 

" iron, 3Fe + 40 = Fe 3 4 ; 

" phosphorus, 2P + 50 = P 2 5 . 

The products are sulphur dioxide, S0 2 , a colorless, 
pungent gas; carbon dioxide, C0 2 , a colorless gas; mag- 
netic oxide of iron, Fe 3 4 , a reddish -brown substance; 
and phosphorus pentoxide, P 2 5 , a white solid which dis- 
solves in water. 

Burning in the Air and Burning in Oxygen. — One can- 
not help noticing a strong resemblance between the 
burning of substances in the air and in oxygen ; and the 
question naturally suggests itself, Are these two acts the 
same in character, or is there a difference between them ? 
To answer this question, we must burn the same sub- 
stances in pure oxygen and in air, and determine whether 
the same products are formed in the two cases, and at the 
same time whether anything else is formed. If we should 
make this comparison in any case, we should find that, 



COMBUSTION. 37 

whether a substance burns in the air or in pure oxygen, 
the same product is formed, and nothing else. It is 
therefore certain that the act of burning in the air is due 
to the presence of oxygen. But substances do not burn 
as readily in the air as in oxygen, and some which burn 
in oxygen do not burn in the air. This is due to the fact 
that only about one-fifth of the volume of the air is oxygen, 
while most of the remaining four-fifths consists of an 
extremely inactive element, nitrogen, which takes little 
part in the process of burning. 

Combustion. — By the term combustion in its broadest 
sense is meant any chemical act that is accompanied by 
an evolution of light and heat. Ordinarily, however, it 
is restricted to the union of substances with oxygen, as 
this union takes place in the air, with evolution of light 
and heat. Substances which have the power to unite with 
oxygen are said to be combustible, and substances which 
have not this power are said to be incombustible. Most 
elements combine with oxygen under proper conditions, 
and are therefore combustible. Most compounds formed 
by the union of oxygen with combustible substances are 
incombustible. For example, the sulphur dioxide, carbon 
dioxide, magnetic oxide of iron, and phosphorus pentoxide, 
formed when sulphur, carbon, iron, and phosphorus are 
burned in oxygen, are incombustible. They contain 
oxygen and they cannot combine directly with more. - 

Kindling Temperature. — We have seen that substances 
do not usually combine with oxygen at ordinary tempera- 
tures, but that in order to effect the union the temperature 
must be raised. If this were not the case, it is plain that 
every combustible substance in nature would burn up, for 
the air probably supplies a sufficient quantity of oxygen 
for this purpose. Some substances need to be heated to a 
high temperature before they will combine with oxygen; 
others require to be heated but little. If we were to 
subject pieces of phosphorus, of sulphur, and of carbon to 
the same gradual rise in temperature, we should find that 
the phosphorus takes fire very easily, only a slight eleva- 



3§ COLLEGE CHEMISTRY. 

tion of temperature being necessary; next in order would 
come the sulphur; and last the carbon. If we were to 
repeat these experiments a number of times, we should 
find that the phosphorus always takes fire at the same 
temperature, and a similar result would be reached in the 
case of sulphur and carbon. Every combustible substance 
has its kindling temperature ; that is, the temperature at 
which it will combine with oxygen. Below this tempera- 
ture it will not combine with oxygen. If a piece of wood 
should be heated to its kindling temperature all at once, 
it would burn up as rapidly as it could secure the necessary 
oxygen; but the burning does not usually take place 
rapidly, for the reason that only a small part of it is at 
any one time heated to the kindling temperature. Watch 
a stick of wood burning, and see how, as we say, ' ' the fire 
creeps " slowly along it. The reason of the slow advance is 
simply this : Only those parts of the stick which are nearest 
the burning part become heated to the kindling tempera- 
ture. They take fire and heat the parts nearest them, and 
so on gradually throughout the length of the stick. 

Slow Oxidation. — Substances may combine slowly with 
oxygen without evolution of light. Thus, if a piece of iron 
is allowed to lie in moist air, it becomes covered with rust. 
The rust is similar to the substance formed when iron is 
burned in oxygen. Both are formed by the union of iron 
and oxygen. Magnesium burns in the air and forms a 
white compound containing oxygen. It burns with in- 
creased brilliancy in oxygen, forming the same compound. 
If left in moist air for some days or weeks, it becomes 
covered with a layer of the same white substance. If this 
is scraped off, and the magnesium again allowed to lie, it 
will again become covered with a layer of the compound 
with oxygen, and this may be continued until the mag- 
nesium has been completely converted into the same sub- 
stance that is formed when it burns in oxygen or in the 
air. Many other cases of slow oxidation might be 
described, some of which, such as the decay of wood, are 
constantly taking place. The most important illustration 



HEAT OF COMBUSTION. 39 

of slow oxidation is that which takes place in our bodies, 
for the food of which we partake undergoes a great many 
changes, some of the substances uniting with oxygen, and 
thus keeping up the temperature of our bodies. This, 
however, is done without evolution of light. We take 
large quantities of oxygen into our lungs in the act of 
breathing. This acts upon various substances presented 
to it, oxidizing them to other forms which can easily be 
got rid of. More will be said in regard to the breathing 
process of animals and plants when the subject of carbon 
and its compounds is taken up. 

Heat of Combustion. — What is the chief difference 
between combustion, as we ordinarily understand it, and 
slow oxidation ? As far as can be judged by a cursory 
examination, it is that in the former there is an evolution 
of light and much heat, while in the latter there is 
apparently but little heat evolved and no light. Kemem- 
bering that the reason why a body gives light is that it is 
heated to a sufficiently high temperature, the problem 
resolves itself into a question of heat. What difference, 
if any, is there between the quantity of heat given off 
when a substance burns, and when it undergoes slow 
oxidation without evolution of light ? Experiment has 
shown that there is no difference. In one case all the heat 
is given off in a short space of time, and therefore the 
temperature of the substance becomes high and it emits 
light. In the other the heat is given off slowly and con- 
tinues for a much longer time, and therefore the tempera- 
ture of the substance does not get high, as surrounding 
substances conduct off the heat nearly as rapidly as it is 
evolved. If, however, we were to measure the quantity 
of the heat, we should find it to be the same in both cases. 

We can measure the heat given off or absorbed in a 
chemical reaction by allowing the reaction to take place 
in a vessel called a calorimeter, so constructed as to pre- 
vent loss of heat, and containing a known weight of water. 
The temperature of the water is noted at the beginning of 
the operation and at the end. A quantity of heat is 



4© COLLEGE CHEMISTRY. 

generally stated by giving the number of grams of water 
which it will raise one degree (Centigrade) in temperature. 
The quantity of heat necessary to raise a gram of water 
from 0° to 1° Centigrade is the unit used in heat measure- 
ments. It is called the calorie. If we say that the 
quantity of heat evolved in any reaction is 250 calories 
(written generally 250 cal.), this means a quantity of heat 
capable of raising the temperature of 250 grams of water 
one degree or of one gram of water 250 degrees in tem- 
perature. Sometimes it is convenient to use a larger unit. 
The quantity of heat required to raise the temperature of 
one kilogram of water one degree serves the purpose. 
This is the large calorie. To distinguish it from the 
smaller one it is written with a capital. Thus, 250 Cal. 
means 250 large calories. The large calorie is obviously 
1000 times greater than the small calorie. 

To repeat, then : By the heat of combustion of a sub- 
stance is meant the quantity of heat given off when a 
certain weight of the substance combines with oxygen. 
In order to avoid confusion it is necessary to have an 
agreement in regard to the weight of substance that shall 
be used as the standard. This may be a gram or any 
other weight, but for the purposes of chemistry it is most 
convenient to take weights in proportion to the combining 
or atomic weights. Thus, by the heat of combustion of 
carbon is meant the quantity of heat evolved in the com- 
bination of 12 grams of carbon with 2 X 16 = 32 grams 
of oxygen. By the heat of combustion of sulphur is meant 
the quantity of heat evolved in the combination of 32.06 
grams of sulphur with 2 X 16 = 32 grams of oxygen, etc. 
Not only is the heat of combustion the same whether 
the union with oxygen takes place slowly or rapidly, but 
the heat evolved in any given chemical reaction is always 
the same, and chemical action is always accompanied by 
an evolution or absorption of heat. 

Heat of Decomposition. — Just as it is true that a definite 
quantity of heat is evolved when two or more elements 
combine chemically, so also it is true that in order to 



CHEMICAL ENERGY AND CHEMICAL WORK. \\ 

overcome the force which holds these elements together 
the same quantity of heat is absorbed. Thus, the heat 
of formation of mercuric oxide, llgO, is 30,060 cal. ; or, 
in other words, when 200.3 grams of metallic mercury and 
10 grams of oxygen combine, 30,000 calories are evolved. 
Now, we have seen that when heat is applied to the com- 
pound it is decomposed into its elements. To effect this 
decomposition, just as much heat is absorbed as was 
evolved in the formation of the compound. 

Chemical Energy and Chemical Work. — Any substance 
which has the power to unite with others can do chemical 
work: it possesses chemical energy. Thus, all combustible 
substances can do work. In uniting with oxygen heat is 
evolved, and this can be transformed into motion. In the 
case of the steam-engine, the source of the motion is the 
burning of the fuel, which is a chemical act. We thus 
see that the source of the power of the steam-engine is 
chemical energy. Substances, on the other hand, which 
have no power to combine with others have no power to 
do chemical work, or they have no chemical energy. So 
far as power to combine with oxygen is concerned, water 
is a substance of this kind, as is also carbon dioxide, the 
gas formed w T hen carbon is burned in oxygen. In order 
that they may do work by combining with oxygen, they 
must first be decomposed, and their constituents put 
together in some form in which they have the power of 
combination. This decomposition of carbon dioxide and 
water is taking place constantly on the earth. All plant- 
life is dependent on it. The products of the action, such 
as the different kinds of wood, have chemical energy, — 
they can do chemical work. This power to do work has 
been acquired from the heat of the sun, which is the main 
force used in decomposing the carbon dioxide and water. 
We have thus a transformation of the sun's heat into 
chemical energy, which is stored up in the combustible 
woods. The quantity of heat which is given off in burn- 
ing wood is believed to be exactly equal to the quantity of 
heat used up in its formation. 



42 COLLEGE CHEMISTRY. 

Oxides. — The compounds of oxygen with other elements 
are called oxides. To distinguish between different oxides, 
the name of the element with which the oxygen is in com- 
bination is prefixed. Thus, the compound of zinc and 
oxygen is called zinc oxide ; that of calcium and oxygen, 
calcium, oxide ; that of silver and oxygen, silver oxide ; etc. 
When an element forms more than one compound with 
oxygen, suffixes are used to distinguish between them. 
Thus in the case of copper there are two oxides which 
have the composition represented by the symbols Cu 2 
and CuO. The former is known as cuprous oxide and the 
latter as cupric oxide. That oxide which contains the 
smaller weight of oxygen in combination with a given 
weight of the other element is designated by the suffix 
ous ; that which contains the larger proportion of oxygen 
is designated by the suffix ic. In other cases the number 
of combining weights of oxygen contained in the com- 
pound is indicated by the name. Thus, manganese dioxide 
is Mn0 2 ; sulphur trioxide is S0 3 ; etc. 

EXPERIMENTS. 
Preparation of Oxygen. 

Experiment 14. — Make some oxygen by heating to redness 10 
to 15 grams (about half an ounce) of manganese dioxide in an 
iron tube closed at one end and connected at the other end by 
means of a cork with a bent glass tube. 

Experiment 15. — Make some oxygen by heating a few grams of 
mercuric oxide in a glass tube closed at one end and connected 
at the other end by means of a cork with a bent glass tube. 

Experiment 16. — Arrange an apparatus as shown in Fig. 3. 
.4 represents a retort of about 100 ccm. capacity. Attach to the 
neck of the retort a piece of rubber tubing B provided with a 
bent glass tip C. Place this tip under the surface of the water 
in D. In A put 4 to 5 grams (about an eighth of an ounce) 
potassium chlorate, and gently heat by means of the lamp. 
Notice carefully what takes place. At first the potassium chlo- 
rate will melt, forming a clear liquid. If the heat is increased, 
the liquid will appear to boil, r>nd it will soon be seen that a gas is 
given off. Now bring the inverted cylinder E filled with water 
over the end of the tube, and let the bubbles of gas rise in the 



MEASUREMENT OE VOLUME OE GASES. 



43 



cylinder. After a considerable quantity of gas has been collected 
in this way the action stops, the mass in the flask becomes solid, 




Fig. 3. 

and apparently the end of the process is reached. But if the 
temperature is raised higher, gas will again begin to come off, and 
in this second stage a larger quantity will be collected than in 
the first. Finally, however, the end is reached, and the sub- 
stance left in the flask remains unchanged, no matter how high 
the temperature may be raised. An examination of the gas col- 
lected will show that a piece of wood will burn in it very readily. 
Explain the changes which have taken place in this experiment. 
Calculate how much oxygen can be obtained by heating 12 grams 
of potassium chlorate. 

Measurement of the Volume of Gases. 

In studying chemical changes it often becomes necessary to 
measure the volume of a gas, and it is important to know what 
precautions must be taken in such cases. For this purpose a 
tube is used which is graduated by marks etched on the outside. 
These marks may either indicate the number of cubic centimetres 
of gas contained in the tube, or the length of the column of gas. 
In the latter case it is of course necessary to determine what vol- 
ume corresponds to a given length of the column. The chief 
difficulty encountered in measuring gas volumes is due to the 
fact that the volume varies with the temperature and pressure. 
When the temperature of a gas is raised one degree Centigrade its 
volume is increased o4 ;T part of its volume at 0°. If, therefore, 
the volume of a gas at 0° is 7, at t° its volume V will be 



7+ w* r > 



or 



7(278 + t) 
273 ' 



44 COLLEGE CHEMISTRY. 

This expression may also be written thus : 

V = V + 0.00366* . V, or V = 7(1 + 0.00366*). 
From these we get the expressions 

Tr 2737' , TT V 

and V = 



273 + t 1 + 0.00366** 

This is called the law of Dalton and Gay-Lussac. 
On lowering the temperature of a gas below 0° the volume is 
diminished at the same rate, that is to say it loses -fa part of its 
volume at 0° for each degree. If this diminution in volume 
should be regular, it is obvious that at a temperature of — 273° 
below 0° the volume of a gaseous substance would be reduced to 
nothing. This temperature is therefore called absolute zero. 
The nearest approach to this temperature thus far produced is 
16° above the absolute 0°. This was produced by Dewar by 
boiling liquid hydrogen. 

It is customary to reduce the observed volume of a gas to the 
volume which it would have at 0°. The correction is easily made 
by the aid of the above formula. Thus, if the volume of a gas is 
found to be 250 cubic centimetres at 15°, and it is required to 
know what the volume would be if the temperature were reduced 
to 0°, the calculation is made thus : In this case the observed 
volume V is 250 cc. ; *, the temperature, is 15°. Substituting 
these values in the equation 

IT 2737 ' M v ' 



273 + *' 1 + 0.00366*' 

we have 

Tr 273 x 250 Tr 250 

V= n „„ , ^ , or V = 



273 + 15 ' 1 + 0.00366 x 15' 

from which we get 236.99 as the value of V. 

But the volume of a gas varies also according to the pressure. 
If the pressure is doubled, the volume is decreased one half ; and 
if the pressure is decreased one half, the volume is doubled, and 
so on. In other words, the volume of a gas varies inversely ac- 
cording to the pressure. Increase the pressure two, three, or four 
times, and the volume becomes one half, one third, or one fourth, 
and vice versa. If the gas has the volume Fat the pressure P, 
and at pressure P' the volume V\ these values are found to 
bear to one another the relations expressed in the equation 
VP = V'P'. 

This is called Boyle's law. 

The pressure is usually stated in millimetres, and reference is to 
the height of a column of mercury which the pressure corresponds 



MEASUREMENT OF VOLUME OF GASES. 



45 



to. A gas contained in an open vessel, or in a vessel over mer- 
cury or water, in which the level of the liquid inside and outside 
the vessel is the same, is under the pressure of the atmosphere. 
What that is we learn from the barometer. As this pressure 
varies, it is necessary to read the barometer whenever a gas is 
measured, and then to reduce the observed volume to certain 
conditions which are accepted as standard. If the gas is meas- 
ured in a tube over mercury or water, and the level of the liquid 
inside the tube is higher than that 
outside, the gas is under diminished 
pressure, the amount of diminution 
depending on the height of the col- 
umn of mercury or water in the 
tube. Thus, if the arrangement is 
as represented in Fig. 4, the height 
of the mercury column above the 
level of the mercury in the trough 
being 100 millimetres, and the pres- 
sure of the atmosphere 760 milli- 
metres of mercury ; then the gas in 
the tube is plainly not under the full 
atmospheric pressure, for the atmos- 
phere is supporting a column of 
mercury 100 millimetres high, and 
the pressure actually brought to bear 
on the gas corresponds to 760 — 100 = 660 mm. of mercury. 
Suppose that in this case the volume of gas actually measured is 
75 cc. Call this V. What would be the actual volume V under 
the standard 760 mm. ? We have seen that 

VP = V'P'. 

Now, in this case P — 760, V = 75, and P' = 660. Therefore, 




Fig. 4. 



760 r 



5 x 660. or V = 



x 660 



■fin 



= 65.13. 



In all cases it is necessary to make a correction similar to this 
in dealing with the volumes of gases. The correction for tem- 
perature and that for pressure may be made in one operation, 
the formula being 



"3V P' 



760(273 + t) 



V = 



V'P' 



760(1 + 0.003660' 



in which V = the volume of the gas at 0° and 760 mm. pressure : 
V = the observed volume; t = the observed temperature; 
P' = the pressure under which the gas is measured. Some of 



4 6 



COLLEGE CHEMISTRY. 



the most important ideas that have been introduced into chem- 
istry with a view to explaining the regularities observed in the 
quantities of substances which act upon 
one another chemically have their origin 
in observations on the conduct of gases. 
It is therefore of the highest importance 
that the student should familiarize himself 
with the meaning of the expression, " the 
volume of a gas under standard condi- 
tions." The presence of water-vapor in a 



by the same amount as it would be in- 
creased if the pressure were diminished by 
a quantity equal to the pressure of the 
water-vapor for the given temperature. 
The formula for making all the correc- 
tions required in determining the volume 
of a gas is 

T AVXP'- a) _ TT _ V'{P'-a) 
760(2734 




or V 



i-\-t) ' " ' 760(1+0.00366)5 

in which the letters V, F', P', and t have 
the same significance as in the last formula 
given, while a is the pressure of water- 
vapor at r. The pressure of water- vapor 
was formerly called "aqueous tension, 11 
aud this term is still often used. The 
amount of the pressure for various tem- 
peratures is to be found from tables. 

A convenient apparatus for measuring 
gas volumes, which simplifies the process, 
is that represented in Fig. 5. It consists 
of two tubes connected at the base by 
means of a piece of rubber tubing, and 
containing water. The tube A is gradu- 
ated, the other is not. The gas the vol- 
ume of which is to be measured is brought 
into the tube .4, with the narrow opening 
at the top, and the other tube is then 
placed at the side of the one containing 
the gas, and its height adjusted so that 
the column of liquid in both tubes is at 
the same level. Under these cireum- 
ances, obviously, the gas is under the atmospheric pressure for 



Fta. 5. 



DETERMINATION OF AMOUNT OF OXYGEN. 



47 



which the necessary correction must of course be made. It is 
also necessary in this case to make the corrections for tempera- 
ture and the pressure of aqueous vapor. It is, further, some- 
times convenient when the gas is measured over water to transfer 
the measuring-tube to a vessel containing enough water to permit 
the immersion of the tube to a point at which the level of the 
liquid inside and outside of the tube is the same. In this case 
the conditions are the same as in the apparatus described in the 
last paragraph. The arrangement is represented in Fig. 6. 



Determination of the Amount of Oxygen liberated 
when a Known Weight of Potassium Chlorate is 
decomposed by Heat. 



Experiment 17. — To determine how much oxygen is given off 
when a known weight of potassium chlorate 
is decomposed by heat, proceed as follows : 
In a small dry hard-glass tube about 10 
cm. long and 5 to 7 mm. internal diameter, 
closed at one end, weigh out on a chemical 
balance about 0.25 gram dry potassium 
chlorate, first weighing the tube empty. 
Introduce just above the potassium chlorate 
a plug of asbestos which has been ignited ; 
then, by means of a blast-lamp, soften the 
upper end of the tube, and draw it out so 
that it has the form shown in Fig. 7. Now 
weigh the tube again. 
Let a = weight of tube empty ; 

I = weight of tube with potassium 

chlorate ; 
c = weight of tube with potassium 
chlorate and plug. 





Fin. 7. A Fig. 0. 

Connect at A by means of a short piece of rubber tubing with 



48 COLLEGE CHEMISTRY. 

the measuring-tube Fig. 5 so that the ends of the two tubes are 
almost in contact with each other, the measuring-tube having 
been previously filled with water to the zero point, and the top 
closed by means of the stop-cock. Open the stop-cock, and now 
heat the potassium chlorate gently at first, and gradually higher 
until no more gas is given off. After the gas has stood for half 
an hour to cool it down to the temperature of the air, adjust 
the two tubes of the measuring apparatus so that the level of 
the water in both is the same ; read off the volume of gas. At 
the same time read the barometer and thermometer ; and now 
make the corrections for pressure and temperature as above di- 
rected. The weight of a litre or 1000 cc. of oxygen at 0° and 760 
mm. pressure is 1.429 grams. Knowing the volume of oxygen 
obtained, calculate the weight of this volume. Kemove the tube 
containing the product left after the decomposition of the potas- 
sium chlorate and weigh it. 

Let d = weight of tube after decomposition of potassium 
chlorate. 

Now b — a = weight of potassium chlorate used ; 
c — d = weight of oxygen given off. 

This weight should, of course, be the same as that calculated 
from the measured volume of oxygen, and should further be in 
agreement with the results calculated by means of the equation 

KC10 3 - KC1 + 30. 

Make all the calculations, and see how nearly the results ob- 
tained agree with what is required by this equation. Should the 
results not be satisfactory the first time, repeat the work. The 
more carefully the work is done, the more nearly will the results 
agree with the equation. 

Experiment 18.— Mix 25 to 30 grams (or about an ounce) of 
potassium chlorate with an equal weight of manganese dioxide 
in a mortar. The substances need not be in the form of powder. 
Heat the mixture in a glass retort, and collect the gas by displace- 
ment of water in appropriate vessels, — cylinders, bell-glasses, 
bottles with wide mouths, etc. It will also be well to collect 
some in a gasometer, such as is commonly found in chemical 
laboratories, the essential features of which are represented in 
Fig. 8. It is made either of metal or of glass. The opening at 
d can be closed by means of a screw cap. In order to fill it 



DFTFRMJNATION OF AMOUNT OF OXYGEN. 



49 



with water, open the stop-cocks and pour the water into the upper 
part of the vessel after hav- 



ing screwed on the cap d. 
When it is full, water will 
flow out of the small tube e. 
Now close all the stop-cocks, 
and remove the cap d. The 
water will stay in the vessel 
for the same reason that it 
will stay in a cylinder in- 
verted with its mouth below 
water. To fill the gasometer 
with gas, put it over a tub 
or sink, and introduce the 
tube from which gas is 
issuing into the opening at 
d. The gas will rise and 
displace the water, which 
will flow out at d. When 
full, put the cap on. To get 
the gas out of the gasometer, 
attach a rubber tube to e, 
pour water into the upper 
part of the gasometer, open Fig. 8. 

the stop-cock a and that at e, when the gas will flow out, and the 
current can be regulated by means of the stop-cock at e. 
The arrangement of the retort is shown in Fig. 9. 





Fig. 9. 



5° 



COLLEGE CHEMISTRY. 



Physical Properties of Oxygen. 

Experiment 19. — Inhale a little of the gas from one of the 
bottles. Has it any taste ? odor ? color ? 




Chemical Properties of Oxygen. 

Experiment 20.— Turn three of the bottles containing oxygen 
with the mouth upward, leaving them covered with glass plates. 
Into one introduce some sulphur in a so- 
called deflagrating-spoon, which is a small 
cup of iron or brass attached to a stout 
wire which passes through a metal 
plate, usually of tin (see Fig. 10). In 
another put a little charcoal (carbon), 
and in a third a piece of phosphorus* 
about the size of a pea. Let them stand 
quietly and notice what changes, if any, 
take place. Sulphur, carbon, and phos- 
phorus are elements, and oxygen is an 
element. It will be noticed that the sul- 
phur and the carbon remain unchanged, 
while some change is taking place in 
the vessel containing the phosphorus, 
as is shown by the appearance of wdrite 
fumes. After some time the phosphorus will disappear entirely, 
the fumes will also disappear, and there will be nothing to 
show us what has become of the phosphorus. If the temperature 
of the room is rather high, it may happen that the phosphorus 
takes fire. If it should, it will burn with an intensely bright 
light. After the burning has stopped, the vessel will be filled 
with white fumes, but these will quickly disappear, and the vessel 
will apparently be empty. What do these experiments prove with 
reference to the action of oxygen on sulphur, carbon, and phos- 
phorus at the ordinary temperature ? 

* Phosphorus should be handled with great care. It is always 
kept under water, usually in the form of sticks. If a small piece is 
wanted, take out a stick with a pair of forceps, and put it under 
water in an evaporating-dish. While it is under the water, cut off a 
piece of the size wanted. Take this out by means of a pair of forceps, 
lay it for a moment on a piece of filter-paper, which will absorb most 
of the water, then quickly put it in the spoon. 



Fig. 10. 



CHEMICAL PROPERTIES OF OXYGEN. 5 1 

Experiment 21. — In a deflagrating-spoon set fire to a little 
sulphur and let it burn in the air. Notice whether it burns 
with ease or with difficulty. Notice the odor of the fumes which 
are given off. Now set fire to another small portion and intro- 
duce it in a spoon into one of the vessels containing oxygen. It 
will be seen that the sulphur burns much more readily in the oxy- 
gen than in the air. Notice the odor of the fumes given off. Is 
it the same as that noticed when the burning takes place in the 
air? 

Experiment 22. — Perform similar experiments with charcoal. 

Experiment 23. — Burn a piece of phosphorus not larger than 
a pea in the air and in oxygen. In the latter case the light 
emitted from the burning phosphorus is so intense that it is pain- 
ful to some eyes. It is best to be cautious. The phenomenon 
is an extremely brilliant one. The walls of the vessel in which 
the burning takes place become covered with a white substance 
which afterwards gradually disappears. 

What differences do you notice between the burning in the air 
and in oxygen ? In the experiments is there any sulphur, or car- 
bon, or phosphorus left behind ? Do the experiments furnish any 
evidence that oxygen takes part in the action ? or that oxygen is 
used up ? 

Experiment 24. — Straighten a steel watch-spring * and fasten 
it in a piece of metal, such as is used for fixing a deflagrating- 
spoon in an upright position ; wind a little thread around the 
lower end, and dip it in melted sulphur. Set fire to this and in- 
sert it into a vessel containing oxygen. For a moment the sul- 
phur will burn as in Experiment 21 ; but soon the steel begins to 
burn brilliantly, and the burning continues as long as there is 
oxygen left in the vessel. Notice that in this case there is no 
flame, but instead very hot particles are given off from the burn- 
ing iron. The phenomenon is of great beauty, especially if ob- 
served in a dark room. The walls of the vessel become covered 
with a dark reddish-brown substance, some of which will also be 
found at the bottom in larger pieces. 

* Old watch-springs can generally be had of any watch maker or 
mender for the asking. A spring can be straightened by unrolling 
it, attaching a weight, and suspending the weight by the spring. 
The spring is then heated up and down to redness with the flame 
of a Bunsen burner. 



52 



COLLEGE CHEMISTRY. 



Oxygen" is used up in Combustion. 

Experiment 25. — Is the odor of the contents of the bottle in 
which the sulphur was burned the same as before the experiment ? 
Iutroduce a stick with a small flame on it successively into the 
vessels used in burning sulphur, carbon, phosphorus, and iron. 
Is oxygen present or not? What evidence have you on this 
point ? 

Experiment 26. — Fill a tube say 30 to 40 cm. (12 to 15 inches) 
long, and 2£ to 3 cm. (1 to 1£ inches) wide, with oxygen, and ar- 
range it in a vessel over water, as shown in Fig. 11. Now fasten 
a small stick of phosphorus to the end of a wire and push it into 
the tube so that about £ to | inch of the phosphorus is above the 
water and exposed to the oxygen. At first no action will take 
place, but after a time w T hite fumes will be 
seen to rise from the phosphorus, and the 
phosphorus will begin to melt. This action 
will be accompanied by a diminution of the 
volume of the oxygen, as will be shown by the 
rise of the w r ater. When the water has risen 
so as to cover the phosphorus, shove the stick 
up so that it is again just above the surface 
of the water. Some of the oxygen will again 
be used up. By working carefully, and re- 
peating this process as many times as may be 
necessary, the oxygen can all be used up with- 
out the active burning of the phosphorus. 
Usually, however, before the action is com- 
pleted, the temperature of the phosphorus be- 
comes so high that it takes fire, when there i s 
a flash of light in the tube and a sudden rise 
of the water, showing that the gas is suddenly 
used up. 

Experiment 27. — Burn a steel watch-spring as directed in 
Experiment 24, with the difference that the spring is passed 
air-tight through a cork which is fitted tightly into the 
neck of the bell-jar (Fig. 12). As the spring burns, the water 
will rise from the vessel in which the bell-jar is standing, and it 
is necessary to pour water into this vessel. When the spring has 
burned near to the cork shove it through so that the burning may 
continue. If the experiment is properly performed the bell-jar 
will be nearly full of water at the end. What does this prove ? 




Fig. 11. 



PRODUCTS OF COMBUSTION. 



S3 



The Products of Combustion weigh more than the 
Substance Burned. 

Experiment 28.— Weigh off about a gram of magnesium ribbon 
in a porcelain crucible. Heat over a Bunseu burner until the 
magnesium has turned to a white sub- 
stance (magnesium oxide). After cool- 
ing, weigh again. Perforin the same ex- 
periment with zinc, tin, and lead. What 
conclusion are you justified in drawing ? 

Experiment 29. — Over each pan of a 
large and rather sensitive balance suspend 
a glass tube filled with pieces of solid 
caustic soda. A balance that will answer 
the purpose very w T ell can be made of 
wood with metal bearings. It may con- 
veniently be about 2£ feet high, with a 
delicate beam about 3 feet long. The best 
tubes for the caustic soda are Argand 
lamp-chimneys, around the bottom of 
which is tied a piece of wire-gauze to pre- 
vent the caustic soda from falling out. 
On one pan of the balance place a candle 
directly under one of the caustic-soda 
tubes, so adjusted that the flame shall be not more than 2£ to 3 
inches below the bottom of the tube. By means of weights placed 
on the other pan establish equilibrium. Now light the candle. 
Slowly, as it burns, the pan upon which it is placed will sink, 
showing that the products of combustion which are partly ab- 
sorbed by the caustic soda are heavier than the candle was. 
While this is by no means an accurate experiment, it is a very 
striking one, and proves beyond question that in the process of 
combustion matter is taken up by the burning body. 




Fig. 12. 



CHAPTEE III. 

A STUDY OF THE ELEMENT HYDROGEN. 

Historical. — Hydrogen was discovered as a distinct sub- 
stance by Cavendish in L766, although it had been 
observed as an inflammable gas before that time. 

Occurrence. It occurs to some extent in the free con- 
dition, and issues from the earth in small quantity in 
some Localities, as, for example, in bhe salt-mines at 
Stassfurt, Germany. It has also been shown to occur in 
enormous quantities in the atmosphere of the sun. On 
bhe earth it occurs chiefly, however, in combination in 
water, of which it forms LI. 19 per cent. It occurs also in 
most substances of animal and vegetable origin, such as 
the various kinds of wood and fruits, and I he tissues of all 
animals. In these products of life if is contained in com- 
bination with carbon and oxygen or with carbon, oxygen, 
and nitrogen. 

Preparation. —The simplest way, theoretically, to pre- 
pare hydrogen is by the decomposition of water by the 
electric current. But (his method is less convenient and 
more expensive than other methods which are available, 
and if is therefore used only under special circumstances. 
If is particularly well adapted to the preparation of small 
quantities of pure hydrogen. 

Some elements, as, for example, sodium and potassium, 
when brought in conflict with water at the ordinary tem- 
perature decompose if and set hydrogen free. When a 
small piece of potassium is thrown upon water, a flame is 
observed at once. When sodium is used, it is seen to form 
a small ball which moves about on the surface of the water 

54 



PREPARATION OP HYbROGEN. 55 

with a hissing sound, bul under ordinary circumstances 
no flame is observed. By applying a flame (<> the ball 
something takes fire and burns. By filling a good-sized 
test-tube with water and inverting it in a larger vessel and 
bringing a sun/1/ piece of sodium wrapped in a piece of 
filter-paper below the mouth of the tube, the sodium will 
rise to the top of the tube when released, and it will then 
be seen that a gas is evolved which gradually depresses the 
water in the tube. A similar experiment with potassium 
gives a similar result. The gas given off in each case is 
hydrogen. By evaporating off the water left in the vessel 
there will be found in each case a white substance of 
marked chemical properties. That formed with the potas- 
sium is known as potassium hydroxide, or caustic potash, 
and has the composition represented by the symbol KOII; 
that formed with the sodium is known as sodium hy- 
droxide, or caustic soda, and is represented by the symbol 
NaOH. The reactions between potassium and sodium 
and water are represented by the equations 

K + H 2 0= KOII +11; and 
Na -f H 2 = NaOH + If. 

Half the hydrogen of the water which is decomposed is 
replaced by the potassium or the sodium, as the case may 
be. These reactions arc partly described by saying that 
the potassium or sodium is substituted for half the 
hydrogen in the water, and the act is called substitution. 
This is a very common kind of chemical action, and we 
shall constantly meet with it in the course of our study. 
The reaction is one of double decomposition or metathesis, 
two substances acting upon each other to form two others. 
Some substances, such as iron, which decompose water 
slowly at the ordinary temperature do so readily at higher 
temperatures. At ordinary temperatures iron decomposes 
water, as is seen in the formation of a coating upon it 
when left in contact with water and the air: At higher 
temperatures when the iron is red-hot it decomposes water 
very readily, and hydrogen can be made in quantity by 



56 COLLEGE CHEMISTRY. 

this means. In the laboratory the iron may be heated in 
a gun-barrel or in a porcelain tube. When steam is passed 
over it the decomposition represented in this equation 
takes place : 

3Fe + 4H 2 = Fe 3 4 + 811. 

The iron combines with the oxygen and liberates the 
hydrogen. 

Carbon, in the form of charcoal or coal, may be used in 
a similar way to effect the decomposition of water and the 
liberation of hydrogen. At a high heat the reaction takes 
place mainly as represented thns : 

C + H 2 = CO + 2H. 

A mixture of two gases, carbon monoxide and hydrogen, 
is thus formed. This mixture is the essential part of the 
gas which has of late years come into such extensive use 
under the name "water-gas." This is formed by passing 
steam over highly-heated anthracite coal or coke. 

By far the most convenient method for making hydrogen 
consists in treating a metal with an acid. Among the 
metals best adapted to the purpose are zinc and iron, and 
indeed zinc is almost exclusively used. As will be seen 
later, acids are substances that contain hydrogen, and are 
characterized by the property that they give up this 
hydrogen easily and take up other elements in the place 
of it. Among the common acids found in every laboratory 
are hydrochloric acid, sulphuric acid, and nitric acid. 
The chemistry of these compounds will be treated of in 
due time; but, as we shall be obliged to use them before 
they are taken up systematically, a feiv words in regard to 
them are desirable in this place. 

Hydrochloric acid is a compound of hydrogen and 
chlorine. It is a gas which dissolves easily in water. It 
is this solution that is used in the laboratory. It is 
manufactured in enormous quantities in connection with 
the manufacture of soda or sodium carbonate. Its chemi- 



PHYSICAL PROPERTIES OF HYDROGEN. 57 

cal Symbol is I TCI. In commerce it is not uncommonly 
called " muriatic acid." 

Sulphuric acid is a compound of sulphur, oxygen, and 
hydrogen in the proportions represented by the formula 
HjS0 4 . It is an oily liquid and is frequently called "oil 
of vitriol." It is manufactured in very large quantities, 
as it plays an important part in many of the most impor- 
tant chemical industries. 

Nitric acid is a compound containing nitrogen, oxygen, 
and hydrogen in the proportions represented by the 
formula HN0 3 . It is a colorless liquid, though, as we get 
it, it is commonly colored straw-yellow. 

When a metal, such as zinc, is brought in contact with 
hydrochloric or sulphuric acid, an evolution of hydrogen 
takes place at once. The reactions are as represented in 
these equations: 

Zn + 2HC1 = ZnCl 2 + 2H; 

Zinc Chloride. 

Zn -{- H 2 S0 4 = ZnS0 4 + 2H. 

Zinc Sulphate. 

Each combining weight of zinc liberates and replaces two 
combining weights of hydrogen. 

The action between iron and these two acids is of the 
same character: 

Fe + 2HC1 = FeCl 2 + 2H; 

Ferrous Chloride. 

Fe + H 2 S0 4 = FeS0 4 + 2H. 

■ Ferrous Sulphate. 

The hydrogen obtained from acids by the action of 
metals is not pure, but it can be purified by treatment 
with appropriate substances. That obtained by the 
decomposition of water by the electric current is pure. 

Physical Properties. — Hydrogen is a colorless, inodor- 
ous, tasteless gas. That made by the action of acids on 
zinc or iron has a somewhat disagreeable odor which is due 
to the presence of other gases in small quantity. It is not 
poisonous, and may therefore be inhaled with impunity. 
We could not, however, live in an atmosphere of hydrogen, 



5 8 COLLEGE CHEMISTRY. 

as we need oxygen. It is the lightest known substance. 
Its specific gravity in terms of the air standard is 0.0696. 
A litre under 760 mm. pressure and at 0° weighs 0.08095 
gram. Under Oxygen it was stated that in chemistry 
hydrogen is commonly taken as the standard of specific 
gravity, and that, hydrogen being unity, the specific 
gravity of oxygen is 15.89. The gas is only slightly solu- 
ble in water. 100 volumes of water take up 1.93 volumes 
of hydrogen. The fact that hydrogen is lighter than the 
air is shown by opening a vessel which contains it and 
turning the mouth of the vessel upward. The gas 
escapes and in a very short time no evidence of its pres- 
ence can be obtained. Light vessels, as, for example, 
soap-bubbles or collodion-balloons, filled with the gas rise 
in the air, and it is used for the purpose of filling large 
balloons. 

At a very low temperature and high pressure, hydrogen 
can be converted into a liquid that boils at — 252°. It 
cannot be liquefied at any temperature above — 242°, no 
matter what pressure it may be subjected to. 

Hydrogen passes readily through porous substances, or 
it diffuses rapidly. This can easily be demonstrated in 
the case of porous earthenware and paper. It also passes 
readily through some metals, as iron and platinum, when 
heated to redness. There is a direct relation between the 
specific gravity of gases and the rate at which they diffuse. 
The lower the specific gravity the more rapid the diffusion. 
The law governing these phenomena is : 

The rate of diffusion of gases is approximately inversely 
proportional to the square roots of their specific gravities. 

The specific gravity of hydrogen being 1 and that of 
oxygen nearly 16 (15.89), the rate of diffusion of oxygen 
is approximately J that of hydrogen. If hydrogen is on 
one side of a porous wall, and oxygen on the other, the 
hydrogen will pass through the wall so much more rapidly 
than the oxygen that there will be an accumulation of 
hydrogen on one side of the wall, and if the vessel were 



CHEMICAL PROPERTIES OE HYDROGEN. 59 

closed there would be increased pressure ou that side. 
The ready passage of gases through porous walls is a 
matter of great importance in connection with the ventila- 
tion of dwellings. Most of the materials used in building 
are porous and permit the passage of gases through them 
in both directions, and change of air is secured in this way 
to some extent. 

Chemical Properties. — Under ordinary circumstances 
hydrogen is not a particularly active element. It does not 
unite with oxygen gas at ordinary temperatures, but, like 
other combustible substances, it must be heated to the 
kindling temperature before it will burn. If a flame is 
applied to it, it takes fire at once. The flame is colorless 
or slightly blue. Generally the flame is somewhat colored 
in consequence of the presence of foreign substances; but 
that it is colorless when the gas is burned alone can be 
shown by burning it as it issues from a platinum tube 
which is itself not chemically acted upon by the heat. 
Although the flame is not luminous it is intensely hot, as 
can be seen by inserting into it a coil of platinum wire, 
which will at once become red-hot and emit light accord- 
ingly. 

The burning of hydrogen in the air, like the burning 
of other combustible substances in the air, consists in a 
union of the gas with oxygen. This has been shown to be 
true by most elaborate experiments on the combustion of 
hydrogen in oxygen and in the air. On the other hand, 
substances which burn in the air are extinguished when 
put in a vessel containing hydrogen. This is equivalent 
to saying that a substance which is uniting with oxygen 
does not continue to unite with oxygen when put in an 
atmosphere of hydrogen, and does not combine with 
hydrogen. The fact is expressed by saying that hydrogen 
does not support combustion. This can be shown by 
holding a vessel filled with hydrogen with the mouth 
downward, and inserting into it a lighted taper supported 
on a wire. The gas takes fire at the mouth of the vessel, 
but the taper is extinguished. 



60 COLLEGE CHEMISTRY. 

Ordinarily we say that hydrogen burns in oxygen, but, 
as the act consists in the union of the two gases, it would 
seem probable that oxygen will burn in an atmosphere of 
hydrogen. This can be shown to be true by a proper 
arrangement of apparatus. If we were surrounded by an 
atmosphere of hydrogen we should speak of oxygen as a 
combustible gas in the same way that we now speak of 
hydrogen as a combustible gas. 

It can easily be shown that, when hydrogen is burned 
either in oxygen or air, water is formed. The simplest 
way to show this is by holding a glass plate or some other 
incombustible object a short distance above a flame of 
hydrogen. It will be seen that drops of water are con- 
densed upon it. 

Hydrogen combines with many other elements besides 
oxygen, and forms some of the most important and inter- 
esting compounds, such as hydrochloric acid, HC1; sul- 
phuretted hydrogen or hydrogen sulphide, H 2 S; ammonia, 
NH 3 ; marsh-gas, CH 4 ; and all the acids. On account of 
its power of combining with oxygen it is used very exten- 
sively in the laboratory for the purpose of extracting 
oxygen from compounds containing it. Thus, when 
hydrogen is passed over heated copper oxide, OuO, it 
combines with the oxygen to form water, and the copper 
is left in the free or uncombined state. The reaction is 
represented thus : 

CuO + 2H = H 2 + Cu. 

A similar reaction takes place when hydrogen is passed 
over highly-heated oxide of iron, Fe 2 3 : 

Fe 2 3 + 6H = 2Fe + 3li 2 0. 

The removal of oxygen from a compound is called 
reduction. Reduction is therefore plainly the opposite of 
oxidation. Any substance that has the power to abstract 
oxygen is spoken of as a reducing agent, just as any sub- 
stance that has the power to add oxygen to a substance, 



COMPARISON OF OXYGEN AND HYDROGEN. 6 1 

or to decompose it by the action of oxygen, is called an 
oxidizing agent. 

A number of metals have the power to absorb a large 
quantity of hydrogen when they are heated in a current of 
the gas. This phenomenon is shown most strikingly by 
palladium, which under the most favorable conditions 
takes up something more than 935 times its own volume 
of hydrogen. The gas i's given up at an elevated tempera- 
ture in a vacuum. When it absorbs hydrogen, palladium 
undergoes marked changes in properties. Its volume is 
increased, and its magnetic and electric properties are also 
changed. It was suggested by Graham that the hydrogen 
held in combination by the palladium is something quite 
different from ordinary hydrogen, and that it must have 
some properties like those of the so-called metals. He 
therefore called the combined hydrogen Uydrogcnium. 
Later studies, especially of liquid hydrogen, have shown that 
there is little in common between hydrogen and the metals. 

Comparison of Oxygen and Hydrogen. — Hydrogen and 
oxygen are different kinds of matter, just as heat and 
electricity are different kinds of energy. Heat can be 
converted into electrical energy, and electrical energy into 
heat, but one element cannot by any means known to us 
be converted into another. They are apparently entirely 
independent of each other. The question will therefore 
suggest itself whether, in spite of their apparent inde. 
pendence, there is not some relation between the different 
elements which reveals itself by similarity in properties. 
It will be found that the elements can be separated into 
groups or families according to their properties. There 
are some elements, for example, which in their chemical 
conduct resemble oxygen markedly. These elements con- 
stitute the oxygen family. So far as hydrogen is con- 
cerned, however, it stands by itself. There is no other 
element which conducts itself like it. If we compare it 
with oxygen, we find very few facts that indicate any 
analogy between the two elements. In their physical 
properties they are, to be sure, similar. Both are color- 



62 COLLEGE CHEMISTRY. 

less, inodorous, tasteless gases. On the other hand, oxygen 
combines readily with a large number of substances with 
which hydrogen does not combine. Oxygen, as we have 
seen, combines easily with carbon, sulphur, phosphorus, 
and iron. It is a difficult matter to get any of these ele- 
ments to combine directly with hydrogen. Further than 
this, substances which combine readily with hydrogen do 
not combine readily with oxygen. The two elements ex- 
hibit opposite chemical properties. What one can do the 
other cannot do. This oppositeness of properties is favor- 
able to combination; for not only do hydrogen and oxygen 
combine with great ease under proper conditions, but, as 
we shall see later, it is a general rule that elements of like 
properties do not readily combine with one another, while 
elements of unlike properties do readily combine. 

EXPERIMENTS. 
Preparation op Hydrogen. 

Experiment 30. — Kepeat Experiment 3 and examine the 
gases. 

Experiment 31 . — Throw a small piece of sodium * on water. 
While it is floating on the surface apply a lighted match to it. 
A yellow flame will appear. This is burning hydrogen, the flame 
being colored yellow by the presence of the sodium, some of which 
also burns. Make the same experiment with potassium. The 
flame appears in this case without the aid of the match. It has 
a violet color, which is due to the burning of some of the potas- 
sium. The gas given off in these experiments is either burned 
at once or escapes into the air. In the case of the potassium it 
takes fire at once, because the action takes place rapidly and the 
heat evolved is sufficient to set Are to it ; in the case of the sodium, 
however, the action takes place more slowly, and the temperature 
does not get'high enough to set fire to the gas. In order to collect 
it unburned, it is only necessary to allow the decomposition to 
take place so that the gas will rise in an inverted vessel filled with 

* The metals sodium and potassium are kept under oil. When a 
small piece is wanted, take out one of the larger pieces from the bottle, 
roughly wipe off the oil with filter-paper, and cut off a piece the size 
needed. It is not advisable to use a piece larger than a small pea. 



PREPARATION OP HYDROGEN. 



63 



water. For this purpose fill a good-sized test-tube with water and 
invert it in a vessel of water. Cut off a piece of sodium not larger 
than a pea, wrap it in a layer or two of filter-paper, and with the 
fingers or a pair of curved forceps bring it quickly below the mouth 
of the test-tube and let go of it. It will rise to the top, the de- 
composition of the water will take place quietly, and the gas 
formed, being unable to escape, will remain in the tube. By 
repeating this operation in the same tube a second portion of gas 
can be made, and so on until the vessel is full. 

Examine the gas and see whether it acts like the hydrogen 
obtained from water by means of the electric current. What evi- 
dence have you that they are the same ? Is this evidence sufficient 
to prove the identity of the two ? 

The metals sodium and potassium disappear in these experi- 
ments, and we get hydrogen. What becomes of the metals ? and 
what is the source of the hydrogen ? If after the action has stopped 
the water is examined, it will be found to contain something in 
solution. It now has a peculiar taste, which we call alkaline ; 
it feels slightly soapy to the touch ; it changes certain vegetable 
colors. If the water is evaporated off, a white substance remains 
behind, which is plainly neither sodium nor potassium. In solid 
form or in very concentrated solution it acts very strongly on 
animal and vegetable substances, disintegrating many of them. 
On account of this action it is known as caustic soda, or, in the 
case of potassium, as caustic potash. 




Fig. 13. 



Experiment 32.— Certain metals that do not decompose water 
at ordinary temperatures, or decompose it slowly, decompose 
it easily at elevated temperatures. This is true of iron. If 
steam is passed through a tube containing pieces of iron heated 



6 4 



COLLEGE CHEMISTRY. 



to redness, decomposition of the water takes place, and the oxy- 
gen is retained by the iron, which enters into combination with 
it, while the hydrogen is liberated. In this experiment an iron 
tube with an internal diameter of from 20 to 25 mm. (about an 
inch) and a gas-furnace are desirable, though a hard-glass tube 
and a charcoal-furnace will answer. The arrangement of the 
apparatus is shown in Fig. 13. 

Experiment 33. — In a cylinder or test-tube put some small 
pieces of zinc, and pour upon it some ordinary hydrochloric acid. 
If the action is brisk, after it has continued for a minute or two 
apply a lighted match to the mouth of the vessel. The gas will 
take fire and burn. If sulphuric acid diluted with five or six 
times its volume of water * is used instead of hydrochloric acid, 





Fig. 14. 



Fig. 15. 



the same result will be reached. The gas evolved is hydrogen. 
For the purpose of collecting the gas the operation is best per- 
formed in a wide-mouthed bottle, in which is fitted a cork with 
two holes (see Fig. 14), or in a bottle with two necks called a 

* If it is desired to dilute ordinary concentrated sulphuric acid 
with water, the acid should he poured slowly into the water while 
the mixture is constantly stirred. If the water is poured into the 
acid, the heat evolved at the places where the two come in contact 
may be so great as to convert the water into steam and cause the 
strong acid to spatter. 






, 



EXPERIMENTS WITH HYDROGEN. 65 

Woulff's flask (see Fig. 15). Through one of the holes a funnel- 
tube passes, and through the other a glass tube bent in a con- 
venient form. 

The zinc used is granulated. It is prepared by melting it in a 
ladle, and pouring the molten metal from an elevation of four or 
five feet into water. The advantage of this form is that it pre- 
sents a large surface to the action of the acids. A handful of 
this zinc is introduced into the bottle, and enough of a cooled 
mixture of sulphuric acid and water (1 volume concentrated acid 
to 6 volumes water) poured upon it to cover it. Usually a brisk 
evolution of gas takes place at once. Wait for two or three 
minutes, and then collect some of the gas by displacement of 
water. When the action becomes slow, add more of the dilute 
acid It will be well to fill several cylinders and bottles with the 
gas, and also a gasometer, from which it can be taken as it is 
needed for experiments. 



Something besides Hydrogen is formed. 

Experiment 34. — After the action is over pour the contents of 
the flask through a filter into an evaporating-dish, and boil off 
the greater part of the water, so that, on cooling, the substance 
contained in solution will be deposited. If the operation is car- 
ried on properly, the substance will be deposited in regular forms 
called crystals. It is zinc sulphate, ZnS0 4 , formed by the sub- 
stitution of zinc for the hydrogen of the sulphuric acid. 

Problems. — How much zinc would it take to give 200 litres of 
hydrogen ? How much zinc sulphate would be formed ? How 
much hydrogen would be formed by the action of 50 grams of 
zinc on sulphuric acid ? How much sulphuric acid would be 
used up ? 



Determination of the Amount of Hydrogen evolved 
when a Known Weight of Zinc is dissolved in 
Sulphuric Acid. • 

Experiment 35. — This determination can be made by means of 
an apparatus such as represented in Fig. 16. The bent tube lead- 
ing from the flask A is drawn out at /?, and a plug of glass-wool 
introduced below the constriction. The other parts of the appa- 
ratus need no description. The flask should have a capacity of 



66 



COLLEGE CHEMISTRY. 



about 40 to 50 cc. ; and the measuring tube C should have a 



" The experiment is conducted in the following manner : D is 
filled with distilled water ; a piece of zinc weighing from 0.150 to 
0.200 gram is placed in the flask ; the pinch-cock E is then opened, 
and the whole apparatus thus filled with water. The apparatus 
is now examined in order to ascertain if gas-bubbles are lodged 
under the stopper F or in the glass-wool. If so, they can usually 
be dislodged without difficulty. If they persist, a few moments' 
boiling of the water in the flask will effect their complete re- 
moval. . . . The eudiometer is now placed over the outlet of the 
delivery-tube, and the greater portion of the water remaining in 
D allowed to flow through the apparatus. Sulphuric acid of the 
concentration ordinarily employed in the laboratory (1 of rLS0 4 
to 4 of HqO) is poured into the reservoir B until it is nearly full. 
The pinch-cock E is then opened, and the water which fills the 




Fig. 16. 

apparatus is displaced by sulphuric acid. The action of the acid 
upon the metal may be facilitated by heat or by adding some 
platinum scraps. When the action is over, the contents of the 
flask are swept through the delivery-tube by again opening the 
pinch-cock E. Finally, the measuring-tube is transferred to a 
cylinder of water, and the volume of the gas read and corrected 
in the usual manner. If hydrochloric instead of sulphuric acid 
has been used, which would be the case when the metal employed 
is aluminium, a little caustic soda should be added to the water 
in the cylinder to which the eudiometer is transferred." 



EXPERIMENTS WITH HYDROGEN. 



67 






A litre of hydrogen at 0° and 760 mm. weighs 0.0895 gram. 
How much does the hydrogen obtained in the experiment weigh ? 
How much ought to have been obtained ? How many cubic cen- 
timetres of hydrogen ought to have been obtained ? 

Try the same experiment, using tin and hydrochloric acid. The 
action takes place as represented in the equation 

Sn + 2HC1 = SnCIa + H 2 . 

It would be well, further, to try the experiment also with iron 
and sulphuric acid, and with aluminium and hydrochloric acid, 
and to calculate from the results the relation between the weights 
of the four metals required to give equal volumes of hydrogen, 
and the volumes of hydrogen given by, say, a gram of each metal. 
The action between iron and sulphuric acid takes place according 
to the equation 

Fe + H 2 S0 4 = FeS0 4 + H 2 . 

That between aluminium and hydrochloric acid is represented by 
this equation : 

Al + 3HC1 = AlCls + 3H. 




Fig. V, 



68 COLLEGE CHEMISTRY. 

Hydrogen is Purified by passing through a Solu- 
tion of Potassium Permanganate. 

Experiment 36. — Pass some of the gas, made by the action of 
zinc on sulphuric acid, through a wash-cylinder containing a solu- 
tion of potassium permanganate ; collect some of it, and notice 
whether it has an odor. The apparatus should be arranged as 
shown in Fig. 17. The solution of potassium permanganate is, 
of course, contained in the small cylinder A, and the tubes so 
arranged that the gas bubbles through it. 
Has the gas any odor or taste or color ? 

Experiment 37. — Place a vessel containing hydrogen with the 
mouth upward and uncovered. In a short time examine the gas 
contained in the vessel, and see whether it is hydrogen. What 
does this experiment prove with reference to the weight of hy- 
drogen as compared with that of the air ? 
Experiment 38. — Gradually bring a vessel containing hydrogen 

with its mouth upward below an 
inverted vessel containing air, in 
the way shown in Fig. 18. After 
the vessel which contained the 
hydrogen has been brought in the 
upright position beneath the 
other, examine the gas in each 
vessel. Which one contains the 
Fig. 18. hydrogen ? 

Experiment 39.— Soap-bubbles filled with hydrogen rise in the 
air. This experiment is best performed by connecting an ordi- 
nary clay pipe by means of a piece of rubber tubing with the 
delivery-tube of a gasometer filled with hydrogen. Small bal- 
loons of collodion are also made for the purpose of showing the 
lightness of hydrogen. 

Hydrogen passes readily through Porous Vessels. 
Diffusion. 

Experiment 40.— Arrange an apparatus as shown in Fig. 19a. 
It consists of a porous earthenware cup, such as is used in gal- 
vanic batteries, fitted with a perforated cork connected with a 
glass tube 2 to 3 feet long. The cork must fit air-tight into the 
mouth of the cup, as well as the tube into the cork. This may 
be secured by shoving the cork into the cup until its outer sur- 




EXPERIMENTS WITH HYDROGEN. 



6 9 




Fig. 196. 



face is even with the edge of the cup, and then covering it care- 
fully with sealing-wax. Put the lower end of the glass tube 
through a cork into one neck of a Woulff's bottle contain- 
ing some water colored with 
litmus or indigo, so that the 
end of the tube is above the 
surface of the water. Through 
the other neck of the bottle 
pass a tube slightly bent out- 
ward and drawn out at the 
end to a fine opening. This 
tube must also be fitted to the 
bottle by an a air-tight cork, 
and its lower end must be 
below the surface of the liquid. 
Now bring a bell-jar (Fig. 196) 
containing dry hydrogen over 
the porous cup, when the liquid 
will be seen to rise in the short, 
bent tube that dips below the 
liquid, and be forced out of it, 
sometimes with considerable 
velocity. Withdraw the bell- 
jar, and bubbles will rise rap- 
idly from the bottom of the 
tube which dips under the 
water, thus showing that air is 
entering the bottle. This is 
due to the diffusion of the 
hydrogen from the porous cup 
into the air. Explain all that you have seen. 





Fig. 19a. 



Fig. 20. 



Chemical Properties oe Hydrogen. 



Experiment 41. — If there is no small platinum tube available, 
roll up a small piece of platinum-foil and melt it into the end of 
a glass tube, as shown in Fig. 20. Connect the burner thus made 
with the gasometer containing hydrogen, and after the gas has 
been allowed to issue from it for a moment, set fire to it. In a 
short time it will be seen that the flame is practically colorless, 
and gives no light That it is hot can be readily shown by hold- 
ing a piece of platinum wire or a piece of some other metal in it. 



7 o 



COLLEGE CHEMISTRY. 



Experiment 42. — Into the flame of burning hydrogen intro- 
duce a small coil of platinum wire. What 
change is observed ? Introduce also a piece 
of magnesium ribbon. Explain the difference 
between the two cases. What becomes of the 
magnesium ? of the platinum ? 

Experiment 43.— Hold a cylinder filled 
with hydrogen with the mouth downward. 
Insert into it a lighted taper held on a bent 
wire, as shown in Fig. 21. The gas takes fire 
at the mouth of the vessel, but the taper is 
extinguished. On withdrawing the taper and 
holding the wick for a moment in the burn- 
ing hydrogen, it will take fire, but on putting 
it back in the hydrogen it will again be extin- 
guished. Other burning substances should 
be tried in a similar way. What conclusions 
are justified by the last two experiments ? 




Pkoduct FOEMED WHEN" Hydeogen is buened. 

Experiment 44. — Hold a clean, dry glass plate a few inches 
above a hydrogen flame. What do you observe ? Kemove what 
is deposited upon the plate, and hold the plate again over the 
flame. Kepeat this a number of times. What does the substance 
deposited upon the plate suggest ? Can you positively say what 
it is ? 



Keduction. 



Experiment 45. — Arrange an apparatus as shown in Fig. 22. 
The flask A contains zinc and dilute sulphuric acid ; the cylinder 
B a solution of potassium permanganate ; the tube G granulated 
calcium chloride ; and the cylinder D concentrated sulphuric 
acid. The object of the potassium permanganate is to purify the 
hydrogen ; the object of the concentrated sulphuric acid and cal- 
cium chloride is to remove moisture from the gas. In the tube E 
put a few pieces of the black oxide of copper, or cupric oxide, 
CuO. After hydrogen has been passing long enough to drive all 
the air out of the apparatus (about two or three minutes if there 
is a brisk evolution) heat the oxide of copper by means of a flame 
applied to the tube. What change in color takes place ? Try the 



EXPERIMENTS WITH HYDROGEN. 



71 



action of nitric acid on the substance before the action and after, 
and note whether there is any difference. What appears in G ? 
Explain what you have seen. 




Fig. 22. 



Experiment 46. — Try the experiment just described, using ferric 
oxide, or oxide of iron, Fe 2 3 , instead of cupric oxide. What is 
the common feature in the two reactions? 



CHAPTER IV. 
STUDY OF THE ACTION OF HYDROGEN ON OXYGEN. 

Burning of Hydrogen. — Attention lias already been 
called to the fact that, when hydrogen burns, water is 
formed. One of the first questions to be answered is, 
What relation exists between the weights of the hydrogen 
burned, the oxygen used up in the burning process, and 
the water formed ? But to weigh gases accurately and to 
collect small quantities of water and weigh it are by no 
means simple operations, and a great deal of work has 
been done upon the problem under consideration. 

Method of Dumas. — The first accurate experiments on 
the combustion of hydrogen are those of Dumas. The 
method employed by this chemist was the following : He 
passed carefully-purified hydrogen over heated copper 
oxide and collected the water formed. The reaction in- 
volved is that represented by the equation 

CuO + 211 = II., + Cu. 

The weight of the oxygen that entered into combination 
with hydrogen was obtained by weighing the vessel con- 
taining the copper oxide before and after the experiment. 
The loss in weight represented the weight of the oxygen 
abstracted. The water was collected by passing the vapor 
into an empty glass vessel in which most of it was con- 
densed, and then through tubes containing molten caustic 
potash and phosphorus pentoxide, substances which have 
a marked power to absorb water and hold it in combina- 
tion. In nineteen experiments he obtained a total amount 
of water weighing 945.430 grams, and the total amount 

72 



MORLEY'S METHOD— EUDIOMETRIC METHOD. 73 

of oxygen tfsed in the formation of this amount of water 
was 840.161 grams. According to these results the ratio 
between hydrogen and oxygen in water expressed in per- 
centages is: 

Oxygen 88.864 

Hydrogen 11.136 



100.000 



Morley's Method. — Morley proceeded as follows: He 
weighed comparatively large volumes of oxygen in the 
form of gas; he then weighed hydrogen after it had been 
absorbed by palladium. The hydrogen being expelled 
from the palladium, it was brought together with the 
oxygen, with which it was caused to combine by means of 
electric sparks. The water thus formed was carefully 
collected in the same apparatus and weighed. Twelve 
experiments gave results that agreed closely with one 
another. The mean of these shows that hydrogen and 
oxgen combine in the ratio of 2 parts of hydrogen to 
15.8792 parts of oxygen. Dumas'' result expressed in the 
same w T ay is 2: 15.961. 

Eudiometric Method. — When hydrogen and oxygen are 
mixed at ordinary temperatures no chemical change takes 
place, and the two gases may be left in contact with each 
other indefinitely without chemical action. If, however, 
a spark is brought in contact with the mixture, violent 
action takes place, accompanied by a flame and explosion. 
The action consists in the sudden combination of the two 
gases to form water. It may be illustrated in a number 
of ways — most simply by filling soap-bubbles with the mix- 
ture of the two gases and applying a flame to them. The 
explosion which ensues is harmless. Plainly, to study the 
combination of hydrogen and oxygen by exploding a mix- 
ture of the gases will require special precautions. It can 
be carried out by the aid of the eudiometer (from evdia, 
good air, and jj.tr pov, a measure, an instrument for deter- 
mining the purity of air). The eudiometer is simply a 
tube graduated in millimetres and having two small plati- 



74 COLLEGE CHEMISTRY. 

num. wires passed through it at the closed end, nearly 
meeting inside and ending in loops outside, as shown in 
Fig. 23. The eudiometer is filled with mercury, inverted 
in a mercury trough, and held in an upright position by 
means of proper clamps. For the purpose of the experi- 
ment a quantity of hydrogen is brought into the tube, and 



Fig. £3. 

its volume accurately measured. About half this volume 
of oxygen is then introduced aud the volume again 
accurately determined, and after the mixture has been 
allowed to stand for a few minutes a spark is passed 
between the wires in the eudiometer by connecting the 
loops with the poles of a small induction coil. Under 
these circumstances the explosion takes place noiselessly 
and with little or no danger. If the interior of the tube 
was dry before the explosion, it will be seen to be moist 
afterwards, and a marked decrease in the volume of the 
gases is also observed. That water is the product of the 
action has been proved beyond any possibility of doubt, 
over and over again. As the liquid water which is formed 
occupies an inappreciable volume as compared with the 
volume of the gases which combine, the decrease in volume 
represents the total volume of the hydrogen and oxygen 
which have combined. Now, if the experiment is per- 
formed with the two gases in different proportions, it will 
be found that only when they are mixed in the proportion 
of 2 volumes of hydrogen to 1 volume of oxygen do they 
completely disappear in the explosion. If there is a larger 
proportion of hydrogen present, the excess is left over; 
and the same is true of the oxygen. It will thus be seen 
that when hydrogen and oxygen combine to form water, 
they do so in the proportion of 2 volumes of hydrogen to 
1 volume of oxygen, or more accurately 2.0027 to 1. 

Calculation of the Results Obtained in Exploding Mix- 
tures of Hydrogen and Oxygen. — Having determined that 



HYDROGEN AND OXYGEN— IVATER-VAPOR. 75 

whenever hydrogen and oxygen combine, they do so in the 
proportion 1 volume oxygen to 2 volumes hydrogen, and 
that when they combine the volume of liquid water formed 
measures so little as to amount to nothing in the measure- 
ments, we know that whenever a mixture of hydrogen and 
oxygen is exploded, no matter in what proportions they 
may be present, the volume of gas which disappears as 
such consisted of 2 volumes of hydrogen and 1 volume of 
oxygen, or, in other words, one-third of the volume which 
disappears was oxygen and two-thirds hydrogen. Take 
this example: A quantity of hydrogen corresponding to 
60 cc. under standard conditions is introduced into a 
eudiometer; 40 cc. of oxygen are added. What contrac- 
tion will there be on exploding the mixture ? Plainly the 
60 cc. of hydrogen will combine with 30 cc. of oxygen. 
The 90 cc. of gas will disappear, and 10 cc. of oxygen will 
remain uncombined. From a total volume of 100 cc, 
therefore, we get a contraction to 10 cc. One-third of the 
contraction represents the oxygen and two-thirds the 
hydrogen. 

Determination of the Volume of Water-vapor formed 
by Union of Definite Volumes of Hydrogen and Oxygen. 
— The experiments which have been described enable us 
to draw the conclusion that hydrogen and oxygen combine 
in certain proportions by volume and by weight, and that 
a definite weight of water is formed; further, that the 
volume of liquid water formed when the two gases combine 
is inappreciable as compared with that of the gases. The 
question remains to be answered, What relation exists 
between the volumes of the combining gases and that of 
the water in the form of vapor ? This can be determined 
by causing the gases to combine in a eudiometer at a tem- 
perature sufficiently high to keep the water in the form of 
vapor. The simplest arrangement for accomplishing this 
is that shown in Fig. 24. A long eudiometer, the upper 
half of which is divided into three equal divisions marked 
on the outside, is filled with mercury and inverted in- a 
bath of mercury. It is then surrounded by a large tube 



7 6 



COLLEGE CHEMISTRY. 



or jacket arranged as shown in Fig. 24. This is connected 
through the cork at the lower end with a vessel from 
which a current of steam can be obtained. The steam is 
passed through the jacket until the 
temperature of the mercury has 
reached that of the steam. A mix- 
ture of hydrogen and oxygen in the 
proportions in which they combine, 
viz., 2 of hydrogen to 1 of oxygen, is 
then introduced into the eudiometer 
so that it is filled to the third mark. 
This must be somewhat above the 
level of the mercury in the bath, so 
that the gases in the eudiometer shall 
be under diminished pressure. On 
now passing the spark the gases unite 
and the water which is formed re- 
mains in the form of vapor, as the 
temperature inside the eudiometer is 
nearly that of boiling water, and the 
vapor is under diminished pressure. 
It is found that the volume of the 
water-vapor is less than that of the 
gases introduced into the eudiometer. In order to secure 
the same pressure as that under which the gases were 
measured, the eudiometer mast be lowered until the 
height of the mercury column in it is the same as it was 
before the explosion. On now measuring the volume of 
water-vapor, it will be found to be two-thirds that occu- 
pied by the uncombined gases. Therefore 2 volumes of 
hydrogen combine with 1 volume of oxygen and form 
2 volumes of water- vapor. It is an interesting fact that 
these simple relations exist between the volumes of the 
combining gases and the volume of the product. We 
shall see that similar relations hold good in the case of 
other gases; and the following general statement is based 
upon a great deal of careful study: 

When two or more gaseous substances combine to form a 




Fig. 24. 



HYDROGEN AND OXYGEN-HEAT. 77 

gaseous compound, (he volumes of the individtial constit- 
uents as well as their sum bear a simple relation to the 
volume of the comjiound. 

This is known as the law of combination by volume. 
As will be seen further on, it has a most important bear- 
ing upon some of the fundamental ideas held in regard to 
the constitution of matter. 

Heat Evolved in the Union of Hydrogen and Oxygen. 
— To get as complete a knowledge as possible of the reac- 
tion which takes place between hydrogen and oxygen we 
have still to determine the amount of heat evolved. The 
heat evolved in burning a gram of hydrogen can be deter- 
mined, and from this we can calculate the heat of formation 
of water which, according to what was said on page 40, is 
the amount of heat evolved by the combination of 2 grams 
of hydrogen with 15.88 grams of oxygen. We are to 
determine the value of x in the equation 

[H 2 , 0] = x cal., 

which expresses in thermochemical language the fact that 
when 2 grams of hydrogen combine with 15.88 grams of 
oxygen x calories of heat are evolved. 

The determination is made by burning a known weight 
of hydrogen in a vessel surrounded by water and arranged 
in such a way that all the heat is absorbed by the water. 
Experiment has shown that when 1 gram of hydrogen is 
burned 34,180 calories are evolved. Or, 

[H 2 , 0] = 68,360 cal. 

No other substance gives as much heat as this in propor- 
tion to the weight used. 

Applications of the Heat formed by the Combination of 
Hydrogen and Oxygen. — To burn hydrogen in the air is, 
as we have seen, a simple matter, but to burn it in oxygen 
requires a special apparatus to prevent the mixing of the 
gases before they reach the end of the tube where the com- 
bustion takes place. The oxyhydrogen blowpipe answers 



78 COLLEGE CHEMISTRY. 

this purpose. This may be constructed in several ways, 
but the simplest is that represented in Fig. 25. It con- 
sists of a tube through which a smaller tube passes. The 




Fig. 25. 



hydrogen is admitted through a and the oxygen through^ b. 
It will be- seen that they come together only at the end of 
the tube. The hydrogen is first passed through and 
lighted; then the oxygen is passed through slowly, the 
pressure being increased until the flame appears thin and 
straight. It gives very little light but is intensely hot. 
Iron wire, steel, copper, zinc, and other metals burn in the 
flame with ease. Platinum vessels are made by melting 
the platinum by means of the oxyhydrogen flame. 

Oxyhydrogen Light. — When the oxyhydrogen flame is 
allowed to play upon some substance which it cannot melt 
or burn, the substance becomes heated so high that it gives 
off an intense light. The substance commonly used for 
this purpose is lime. Hence the light is often called the 
iime -light. It is also known as the Drummond light. 

The hydrogen is first allowed to pass through the stop- 
cock, and lighted, when the oxygen is admitted. The 
flame plays against the piece of lime, and from this the 
light is given off when it has acquired a high temperature. 
Coal-gas may be and is generally used instead of hydrogen. 
Cylinders of compressed coal-gas and of oxygen can be 
bought, and the gases so prepared are used for the purpose 
of projecting pictures upon screens in illustrating lectures 
and for other similar purposes. 

Summary. — In our study of the action of hydrogen on 
oxygen, we have learned: (1) the relations between the 
weights of the two gases which act upon each other; (2) 
the relations between the volumes of the combining gases; 



COMPOSITION OF WATER. 79 

(3) the relations between the volumes of the combining 
gases and that of the water- vapor formed; and (4) the 
amount of heat evolved when a given weight of hydrogen 
combines with oxygen. It remains for us to study more 
carefully the product formed. This is water. 



EXPERIMENTS. 

Composition of Water. 

^Experiment 47. — Arrange the apparatus shown in Fig. 22 
with a straight tube instead of the bent tube E, and connect this 
with a small bent tube containing calcium chloride, as shown in 
Fig. 26. Weigh tube E empty, and after the cupric oxide has 
been put into it. This gives the weight of the cupric oxide. 
Weigh the tube F before the experiment. Now proceed as in 
Experiment 45. In this case all the water formed by the action 
of the hydrogen on the cupric oxide will be absorbed by the cal- 
cium chloride in tube F. This tube will therefore gain in weight, 
and as oxygen is removed from the cupric oxide, tube E will 




Fig. 26. 

lose in weight. After the reduction is complete weigh tube E 
and tube F again. 

Let x = weight of tube E + cupric oxide before the experi- 
ment ; 
y = weight of tube E + copper after the experiment. 
Then x — y = weight of oxygen removed from the cupric oxide. 

Let a = weight of tube F before the experiment, 
and b = " *> " after " " 

Then b — a = weight of water formed. 

If the experiment is properly performed, it will be found that 

the ratio y is very nearlv -z. Or the result may be stated 

o — a 9 

thus : In nine parts of water there are eight parts of oxygen. 



8o 



COLLEGE CHEMISTRY. 



Experiment 48. — The tubes in the apparatus used in Experi- 
ment 3, or some other similar apparatus, should be graduated. 
Let the gases formed by the action of the electric current, as in 
Experiment 3, rise in the tubes, and observe the volumes. It 
will be seen that when one tube is just full of gas, the other, if 
it is of the same size, will be only half full. On examining the 
gases the larger volume will be found to be hydrogen, and the 
smaller volume oxygen. What are the relative weights of equal 
volumes of hydrogen and oxygen ? In what proportion by weight 
are the two gases obtained from water in this experiment ? How 
does this result agree with that obtained in the preceding experi- 
ment ? Does this experiment prove that water consists only of 
hydrogen and oxygen ? 

Experiment 49.— Pass hydrogen from a generating-flask or a 
gasometer through a tube containing some substance that will 
absorb moisture, as all gases made in the ordinary way and 
collected over water are charged with moisture. The calcium 
chloride should be in granulated form, not powdered. After 
passing the hydrogen through the calcium chloride, pass it 
through a tube ending in a narrow opening, and set fire to it. If 
now a dry vessel is held over the flame, drops of water will con- 
dense on its surface and run down. A convenient arrangement 
of the apparatus is shown in Fig. 27. 




Fig. 21 



A is the calcium chloride tube. Before lighting the jet hold a 
glass plate in the escaping gas, and see whether water is deposited 
on it. Light the jet before putting it under the bell-jar, other- 



EUDIOMETRIC EXPERIMENTS. Si 

wise if hydrogen is allowed to escape into the vessel it will con- 
tain a mixture of air and hydrogen, and this mixture, as we shall 
soon see, is explosive. 

Experiment 50. — Mix hydrogen and oxygen in the proportion 
of about 2 volumes of hydrogen to 1 volume of oxygen in a gas- 
ometer. Fill soap-bubbles, made as directed in Experiment 30, 
with this mixture and allow them to rise in the air. As they rise, 
bring a lighted taper in contact with them, when a sharp explo- 
sion will occur. Great care must be taken to keep all flames 
away from the vicinity of the gasometer while the mixture is in 
it. This experiment is conveniently performed by hanging up, 
about six to eight feet above the experiment-table, a good-sized 
tin funnel-shaped vessel with the mouth downward. Now place 
a gas jet or a small flame of any kind at the mouth of the vessel. 
If the soap-bubbles are allowed to rise below this apparatus 
they will come in contact with the flame and explode at once.* 

What does this experiment show ? Does it give any informa- 
tion in regard to the composition of water ? 

Eudiometric Experiments. 

Experiment 51. — The general method of studying the combin- 
ation of hydrogen and oxygen by means of the eudiometer was 
described in the text (see p. 73). To what was there said it need 
only be added that, in exploding the mixture in the eudiometer, 
the latter should be held down firmly, by means of a clamp, 
against a thick piece of rubber cloth placed on the bottom of the 
mercury-trough. In making the measurements of the volume of 
the gases and the height of the mercury column, care must be 
taken to have the eudiometer in a perpendicular position. This 
can be secured by means of plumb-lines suspended from the ceil- 
ing and reaching nearly to the table, by which the position of the 
eudiometer can be adjusted. 

OXYHYDROGEN BLOWPIPE. 

Experiment 52. — Hold in the flame of the oxyhydrogen blow- 
pipe successively a piece of iron wire, a piece of a steel watch- 
spring, a piece of copper wire, a piece of zinc, a piece of platinum 
wire. 

* The same apparatus may be used in experimenting with soap- 
bubbles filled with hydrogen. 



82 COLLEGE CHEMISTRY. 

Experiment 53. — Cut a piece of lime of convenient size and 
shape, say an inch long by three-quarters of an inch wide, and the 
same thickness. Fix it in position so that the flame of the oxy- 
hydrogeu blowpipe will play upon it. The light is very bright, 
but by no means as intense as the electric light. 



CHAPTER V. 
WATER. 

Historical. — Water was long thought to be an elementary 
substance until, towards the end of the last century, the 
discovery of hydrogen and oxygen, and of the nature of 
combustion, led to the discovery of its composition. 

Occurrence. — Besides the form in which water occurs in 
such enormous quantities in the earth, it also occurs in 
forms and conditions which prevent its immediate recogni- 
tion. Thus all living things contain a large proportion of 
water, which can be driven off by heat. If a piece of wood 
or of meat is heated, liquids pass off, and by purification 
these can be shown to consist mainly of water. The pro- 
portion of water in animal and vegetable substances is very 
great. If the body of a man weighing 150 pounds were 
to be put in an oven and thoroughly dried, there would be 
only about 40 pounds of solid matter left, most of the rest 
being water. 

Water also occurs in another form in which it does not 
directly reveal its presence. This is as water of crystal- 
lization. Many natural and manufactured chemical com- 
pounds give off water when heated. If, for example, the 
zinc sulphate formed in the preparation of hydrogen from 
zinc and sulphuric acid is dried by exposure to the air or 
by pressing between layers of filter-paper, it will be found 
that, when heated in a dry tube, it gives off water, and at 
the same time changes its appearance. The same is true 
of gypsum which is found in nature, and of copper sul- 
phate or blue vitriol. In this last case the loss of water is 



84 COLLEGE CHEMISTRY. 

accompanied by a loss of color. After all the water is 
driven off, the powder left behind is white. 

Man} 7 compounds when deposited from solutions in water 
in the form of crystals combine with definite quantities of 
water. This water is not present as such, but is held in 
chemical combination. Hence the substance does not 
appear moist, though it may contain more than half its 
weight of water. This water of crystallization is, in some 
way which is not understood, essential to the form of the 
crystals. If it is driven off, the crystals generally crumble 
to pieces. Some compounds combine with different quan- 
tities of water under different circumstances, the form of 
the crystals varying with the quantity of water held in 
combination. 

Compounds differ greatly as regards the ease with which 
they give up water of crystallization. In general, it is 
given off when the compound containing it is heated to the 
temperature of boiling water. But some compounds give 
it up at the ordinary temperature. This is true of 
sodium sulphate or Glauber's salt, which contains a quan- 
tity of water of crystallization represented by the formula 
Na 2 SO 4 .10H 2 O. If some of the crystals are allowed to lie 
exposed to the air they undergo a marked change in the 
course of an hour or two. They lose their lustre and 
gradually crumble to pieces. Substances which lose their 
water of crystallization by exposure to the air are said to 
be efflorescent. 

Some compounds if deprived of their water of crystal- 
lization will take it up again when allowed to lie in an 
atmosphere containing moisture, as, for example, in the 
air. It is well shown by the compound calcium chloride, 
Ca01 2 . If exposed to the air it will be noticed that it soon 
has a moist appearance, and if it is allowed to lie long 
enough it will dissolve in the water which it absorbs from 
the air. Substances which absorb enough water from the 
air to cause them to dissolve are said to be deliquescent. 

Formation of Water and Proofs of its Composition — If 
we had not already learned, in studying the action of 



FORMATION OF WATER. 85 

hydrogen upon oxygon, that water is composed of these 
two gases, we should first subject it to analysis. For this 
purpose Ave should have to bring it under the influence of 
a number of reagents and study its conduct. If we should 
pass an electric current through it in the proper way, we 
should observe that a gas rises from each pole. By placing 
each pole under the mouth of an inverted tube filled with 
water the gases are easily collected. When one of the 
tubes has become full of gas the other one will be only half 
full. An examination will show that the gas in the tube 
which is filled is hydrogen, whereas that in the one which 
is half filled is oxygen. By decomposition of water by 
means of the electric current, then, there are obtained two 
volumes of hydrogen for each volume of oxygen. We 
already know the relative weights of equal volumes of the 
two gases, so that we can easily calculate the relative 
weights of the gases obtained in the experiment. The 
ratio of the weights of equal volumes of hydrogen and 
oxygen is 1 : 15.89. Therefore, if we have 2 volumes of 
hydrogen combined with 1 volume of oxygen, the ratio 
between the weights is 2.0027 : 15.89 or 1 : 7.95/ Although 
we know from the experiment referred to that hydrogen 
and oxygen are obtained from water in certain proportions, 
it does not follow that this is the composition of water. 
For it may be that other elements besides hydrogen and 
oxygen are contained in it, and it may be also that all the 
hydrogen and oxygen are not set free by the action of the 
electric current. We might determine whether either of 
these possibilities is true or not by decomposing a weighed 
quantity of water, and weighing the hydrogen and oxygen 
obtained from it. If we should find that the sum of the 
weights of hydrogen and oxygen is equal to the weight of 
the water decomposed, this fact would be evidence that 
only hydrogen and oxygen are contained in water, and 
that they are present in the proportions stated. The same 
thing can be satisfactorily proved by causing hydrogen and 
oxygen to combine, or by effecting the synthesis of water. 
How this can be done has already been pointed out. It 



86 COLLEGE CHEMISTRY. 

was shown, in the first place, that by burning hydrogen 
in oxygen water is formed. This proves that water con- 
sists of hydrogen and oxygen, but it does not furnish any 
proof as to the relation between the weights of the gases 
that combine. It is a qualitative synthesis. Other 
methods were described, the object of which was to show 
in what proportion by weight and by volume hydrogen and 
oxygen combine to form water. These methods are ex- 
amples of quantitative syntheses. The results proved that 
to form water hydrogen and oxygen combine in the pro- 
portion of 1 volume of oxygen to 2 of hydrogen; and it 
therefore follows that the decomposition of water which is 
effected by the electric current is complete. 

Properties of Water. — Pure water is tasteless and in- 
odorous, and in small quantities colorless. Thick layers 
are, however, blue. This is seen by filling a long tube 
with carefully purified water, and examining it by trans- 
mitted light, when it appears blue. Some mountain lakes 
also have a marked blue color. When cooled, water con- 
tracts until it reaches the temperature of 4°. At this 
point it has its maximum density. If cooled below this it 
expands, and the specific gravity of ice is somewhat less 
than that of water. Hence ice floats on water. If this 
were not so, in cold climates the water in the streams 
would freeze solid. As it is, the lower layers of water are 
protected by the ice and the cold water just below it, which 
are poor conductors of heat. Water can be cooled down 
below its freezing temperature, or 0°, if it is kept perfectly 
quiet, protected from the air, or cooled in capillary tubes. 
Water thus cooled down will suddenly solidify when dis- 
turbed, and then its temperature rises to 0°. Water boils 
at 100° under 760 mm. pressure. Increased pressure 
raises the boiling-point, and decreased pressure lowers it. 

Chemical Properties of Water. — Water is a very stable 
chemical compound. An indication of the great energy 
of the act of combination of hydrogen and oxygen is given 
in the great evolution of heat. In order to decompose 
water by heat, as much heat must be added as is evolved 



CHEMICAL PROPERTIES OE WATER. 87 

when it is formed. At high temperatures it is decomposed 
into its elements. The decomposition begins at 1000° and 
is half complete at 2500°. On cooling the mixed gases 
they recombine. The amount of decomposition is constant 
for a given temperature. This kind of gradual decom- 
position of a compound by heat, followed by a recombina- 
tion on reversing the conditions, is called dissociation. It 
is a common phenomenon in chemistry, and farther on we 
shall have occasion to study it more fully. Water is 
decomposed by an electric current, when something is 
added to it to make it a conductor of electricity, and partly 
by contact with sodium and potassium at ordinary tem- 
peratures, and by iron and carbon at higher temperatures. 
It combines directly with a large number of substances in 
the form of water of crystallization, and with others to 
form definite chemical compounds called hydrates or 
hydroxides. Thus the oxides of potassium, K 2 0, and of 
sodium, Na 2 0, combine with water with evolution of much 
heat to form the compounds potassium hydroxide, KOH, 
and sodium hydroxide, NaOH, which, it will be remem- 
bered, are also formed by the action of the elements potas- 
sium and sodium on water with liberation of hydrogen. 
The reactions between the oxides and water are represented 
by the equations 

K 2 + H 2 = 2KOH; 
Na 2 + H 2 = 2NaOH. 

Similarly, lime or calcium oxide, CaO, acts upon water 
with evolution of heat, as is observed in the process of 
slaking. The change is like that which takes place with 
potassium and sodium oxides, and is represented thus : 

CaO -f H 2 = Ca0 2 H 2 [or Ca(OH)J. 

The product represented by the symbol Ca(OH) 2 is known 
as calcium hydroxide or slaked lime. In the same way 
barium oxide, BaO, forms barium hydroxide : 

BaO + H 2 = Ba0 2 II 2 [or Ba(OII)J. 



88 COLLEGE CHEMISTRY. 

We shall meet with many examples of this kind of action 
in our study of chemical reactions, and we shall' see that 
the hydroxides form two of the most important classes of 
compounds, known as acids and bases. The hydroxides 
of potassium, sodium, calcium, and barium are, for exam- 
ple, bases; while certain hydroxides containing sulphur, 
nitrogen, and carbon are acids, — such as sulphuric acid, 
S0 2 (OH) 2 , nitric acid, N0 2 (OH), and carbonic acid, 
CO(OH) 9 . It is not believed that water as such is con- 
tained in these hydroxides. Nevertheless, when heated 
many of them give oh* water. Thus, when heated to a red 
heat, calcium hydroxide is decomposed into the oxide and 
water according to the equation 

Ca(OH) 2 = CaO + H 2 0. 

Many substances that contain hydrogen and oxygen act 
in the same way. This is due to the great stability of 
water even at elevated temperatures. As the temperature 
becomes higher and higher the constituents of the com- 
pound are held together less and less firmly. When a 
point is reached at which the attraction of the hydrogen 
for the oxygen is greater than that required to hold the con- 
stituents together, a rearrangement takes place, and com- 
pounds which are stable at the higher temperature are 
formed. 

Water as a Solvent. — With a great many substances 
water forms unstable compounds, the nature of which 
cannot at present be explained. These unstable com- 
pounds are called solutions. It is known that many solids, 
liquids, and gases when brought into water disappear and 
form colorless liquids which look like water. Some give 
colored liquids of the same color as the substance dissolved, 
and others give liquids which have a color quite different 
from that of the substance dissolved, On the other hand, 
there are many substances which do not form such com- 
pounds with water or which, as we say, are insoluble in 
water. In a solution the particles of the substance dis- 



WATER AS A SOLVENT. 89 

solved are in some way held in combination by the particles 
of the liquid. If a very small quantity of substance is dis- 
solved in a large quantity of water and the solution 
thoroughly stirred, the dissolved substance is uniformly 
distributed throughout the liquid, as can be shown by 
refined chemical methods. That the dissolved substance 
is everywhere present in the solution can be shown further 
by the aid of certain dye-stuffs, as, for example, magenta. 
A drop of a concentrated solution of the substance brought 
into many gallons of water imparts a distinct color to all 
parts of the liquid. An experiment of this kind gives 
some idea of the extent to which the division of matter 
can be carried. For it is evident that in each drop of the 
dilute solution there must be contained some of the color- 
ing matter, though the quantity must be what we should 
ordinarily speak of as infinitesimal. While there seems to 
be no limit to the extent to which a solution can be diluted 
and still retain the dissolved substance uniformly dis- 
tributed through its mass, there is a limit to the amount 
of every substance that can be brought into solution, and 
this varies with the temperature, and, in the case of gases, 
with the pressure. Some substances are easily soluble, 
others are difficultly soluble. When the solutions are boiled 
the water simply passes on 3 and leaves the dissolved sub- 
stance behind, if it is a non-volatile solid. If, however, 
the substance in solution is a liquid, a partial separation 
will take place, the extent of the separation depending 
largely upon the difference between the boiling-points of 
the water and the other liquid. A complete separation of 
two liquids by boiling is difficult and in most cases im- 
possible. If, finally, the substance in solution is a gas, it 
generally passes off when the solution is heated, though in 
some cases water is given off leaving the gas in solution, 
which of course then becomes more concentrated. When 
a certain concentration is reached a solution of the gas 
passes over. It is probable that in these cases the gas is 
in a condition of chemical combination with the water. 
Solutions in general seem to differ from true chemical 



90 COLLEGE CHEMISTRY. 

compounds in some important particulars, and also from 
mere mechanical mixtures. Definiteness of composition 
appears to be characteristic of chemical compounds, or, at 
least, it is characteristic of a large number of compounds 
which we call chemical compounds. But solutions have 
no definite composition. We can dissolve any quantity of 
a substance from the minutest particle to a certain fixed 
quantity, and the solutions formed are uniform and appear 
to be just as truly solutions as that which contains the 
largest quantity which can be held in combination. On 
the other hand, in a mere mechanical mixture the constit- 
uents may be present in all proportions, while this is not 
true of solutions. The subject of solution is under inves- 
tigation, and it has been made highly probable that some 
substances, when dissolved, are partly or wholly broken 
down into smaller parts charged with positive and negative 
electricity respectively. These smaller parts are called 
ions. 

Solution as an Aid to Chemical Action. — When it is 
desired to secure the chemical action of one solid substance 
upon another, it is generally necessary to bring them 
together in solution. One reason why they do not act 
readily when mixed in the solid condition is to be found 
in the fact that, under these circumstances, their particles 
remain separated by sensible distances, no. matter how 
finely the mixture may be powdered. If, however, the 
substances are dissolved, and the solutions poured together, 
the particles of the liquid move so freely among one another 
that they come in intimate contact, thus facilitating 
chemical action. Many substances which do not act upon 
one another at all when brought together in dry condition 
act readily when brought together in solution. It is 
believed that this is due principally to the splitting of the 
compounds into ions by the action of the water. These 
ions are free to act upon other ions which may be brought 
into the same solution. This idea will be developed 
farther on in other connections. Although it is highly 
probable, then, that when a reaction takes place in a water 



NATURAL WATERS. 9 1 

solution the water itself plays a very important part, the 
reaction is generally represented by an equation in which 
the water does not appear. Of course, such an equation 
is imperfect, hut it answers certain purposes quite satisfac- 
torily, and may be used without danger of confusion. 
Thus, when hydrochloric acid acts upon zinc, hydrogen is 
liberated and zinc chloride is formed. What we call 
hydrochloric acid in the laboratory is the liquid which is 
formed by the absorption of hydrochloric acid gas, HOI, 
by water. When this liquid is used, however, the chem- 
ical act which makes itself known to us is that which gives 
hydrogen gas and zinc chloride in solution. Apparently 
this reaction is independent of the water, and it may be 
represented thus : 

Zn + 2HC1 = ZnC] 2 -f 2H. 

Sometimes such reactions are written as follows in order 
to express the fact that water is present, though, a's will 
be observed, no attempt is made to tell what part the 
water plays: 

Zn + 2HC1 + Aq = ZnCl 2 + 2H + Aq; or 
Zn + 2HC1 + H 2 = ZnCl 2 + 2H + H 2 0. 

There is no objection to this, certainly, but it is question- 
able whether, considering the purposes for which chemical 
equations are generally used, this increases their value. 

Natural Waters. — All water found in nature is more or 
less impure or, in other words, contains something in 
solution. In the first place, waters which are exposed to 
the air dissolve some of the gases of which the air is com- 
posed, as oxygen, nitrogen, and carbon dioxide. Again, 
natural waters necessarily are in contact with the earth; 
they always dissolve some of the earthy substances; and, 
finally, many waters come in contact with animal and 
vegetable substances and dissolve something from these. 
The water which is carried up as vapor from the surfaces 
of natural bodies of water is approximately pure. When 
this is precipitated as rain it dissolves certain substances 



92 COLLEGE CHEMISTRY. 

from the air, and the first rain that falls during a storm is 
always more or less contaminated. In a short time, how- 
ever, the air becomes washed and the rain which falls 
thereafter is approximately pure water. If it remains in 
contact with insoluble rocks, as, for example, quartzite or 
sandstone, it remains pure, and mountain-streams which 
flow over sandstone beds are, in general, the purest. 
Water which flows over limestone dissolves some of this 
and becomes "hard." A similar change is brought about 
in water by contact with gypsum and magnesium sulphate. 
The condition of hardness will be taken up more fully 
under calcium and magnesium compounds. The many 
varieties of mineral springs have their origin in the pres- 
ence in the earth of certain substances which are soluble 
in water. Among those most frequently met with in 
solution in natural waters are carbonic acid, sodium car- 
bonate, sodium sulphate or Glauber's salt, sodium chloride 
or common salt, magnesium sulphate, iron carbonate, and 
sulphuretted hydrogen. Effervescent waters are those 
which contain a large quantity of carbonic acid in solution 
and give off carbon dioxide gas when exposed to the air. 
Clialybeate waters are those which contain some compound 
of iron in solution; sulphur waters contain the gas, sul- 
phuretted hydrogen. Common salt occurs in large 
quantities in different parts of the earth. As it is easily 
soluble in water, many streams contain it; and as most 
streams find their way to the ocean, we see one reason 
why the water of the ocean should be salt. 

As streams approach the habitation of man they are 
subjected to a serious cause of contamination. The 
drainage from the neighborhood of human dwellings is 
very apt to find its way into a near stream. The sub- 
stances thus carried into the stream undergo decomposition 
and give rise to the formation of larger or smaller quanti- 
ties of new products, some of which have the power to 
produce disturbances when taken into the system, and 
others to produce disease. This condition of things is 
most strikingly illustrated by the case of a large town 



PURIFICATION OF WATER. 93 

situated on the banks of a river. It frequently happens 
that the water of the river is used for drinking purposes, 
and it also frequently happens that the water is contami- 
nated by drainage. Kiver water when once contaminated 
by drainage tends to become pure again by contact with 
the air, the change consisting largely in the slow oxidation 
of the substances which are of animal or vegetable origin, 
and their conversion into harmless products. If water is 
to be used for drinking purposes, however, it is not well 
to rely too much upon this process of purification. So 
much has of late years been said about drinking-water that 
excessive alarm has been created, and water is no doubt 
frequently held responsible for sickness with which it has 
nothing to do. In some places the war against the water- 
supply has been carried so far that those who can afford 
it drink only artificially purified and distilled water. It 
is undoubtedly well to be cautious, but it is possible to be 
too cautious. 

What Constitutes a Bad Drinking-water. — A good 
drinking-water should be free from odor and taste and 
should not contain anything that can act injuriously upon 
the system. It is, however, difficult to decide by chemical 
means whether the water contains anything injurious or 
not, as there may be a very minute quantity of an ex- 
tremely injurious substance, for example a disease germ, 
present, and chemical analysis would be powerless to detect 
it. On the other hand, water which is very considerably 
contaminated by sewage may be harmless, and yet the 
latter might be pronounced "bad" and the former 
" good." The rule generally adopted by chemists in deal- 
ing with water is to pronounce any water dangerous that 
is contaminated by sewage. Such contamination can 
generally be detected by analysis or by analysis and inspec- 
tion of the sources. 

Purification of Water. — Impure water may become 
purer by natural methods as has been stated, and it may 
be rendered fit for drinking purposes by filtering through 
such substances as charcoal, sand, spongy iron, etc. A 



94 COLLEGE CHEMISTRY. 

filter, no matter of what it may be made, will not, how- 
ever, remain efficient for any length of time, as the sub- 
stances contained in the impure water are retained by it 
and, after a time, it becomes a source of pollution instead 
of a purifier. For refined work in chemistry pure water 
is prepared by distilling natural waters. The process of 
distillation consists in boiling the water and then passing 
the steam through a tube or system of tubes surrounded 
by cold water. Thus the steam is condensed, and the dis- 
tilled water is approximately pure. Of course, it is neces- 
sary that the tubes in which the condensation takes place 
should be of such material that water does not act upon it 
to any extent. The materials used are glass, block tin, 
and platinum. Chemically pure water is a very rare sub- 
stance even in the best chemical laboratories. The slight 
impurities present in ordinary distilled water are not, 
however, of special importance under ordinary circum- 
stances. 

EXPEKIMENTS. 

Organic Substances contain Water. 

Experiment 54.— In dry test-tubes heat gently various organic 
substances as a piece of wood, fresh meat, fruits, vegetables, etc. 

Water of Crystallization. 

Experiment 55. — Take some of the crystals of zinc sulphate 
obtained in Experiment 34. Spread them out on a layer of filter- 
paper, and finally press two or three of them between folds of the 
paper. Examine them carefully. They appear to be quite dry, 
and in the ordinary sense they are dry. Put them in a dry tube 
and heat them gently, when it will be observed that water con- 
denses in the upper part of the tube, while the crystals lose their 
lustre, becoming white and opaque, and at last crumbling to 
powder. 

Experiment 56. — Perform a similar experiment with some 
gypsum, which is the natural substance from which " plaster of 
Paris " is made. 



PURIFICATION OF WATER. 



95 



Experiment 57. — Heat a few small crystals of copper sulphate, 
or blue vitriol. In this case the loss of water is accompanied by 
a loss of color. After all the water is driven off, the powder left 
behind is white. On dissolving it in water, however, the solution 
will be seen to be blue ; and if the solution is evaporated until 
the substance is deposited, it will appear in the form of blue 
crystals. 

Efflorescent Salts. 

Experiment 58. — Select a few crystals of sodium sulphate 
which have not lost their lustre. Put them on a watch-glass, 
and let them lie exposed to the air for an hour or two. They 
soon lose their lustre, and undergo the changes noticed in heat- 
ing zinc sulphate. 

Deliquescent Salts. 

Experiment 59.— Expose a few pieces of calcium chloride to 
the air. Its surface will soon give evidence of the presence of 
moisture, and after a time the substance will dissolve in the 
water which is absorbed. 

Purification of Water by Distillation. 
Experiment 60. — In an apparatus like that shown in Fig. 28 




Fig 



distil a dilute solution of copper sulphate or some other colored 
substance. A slow current of cold water must be kept running 



96 COLLEGE CHEMISTRY. 

through the condenser by connecting the lower rubber tube with 
a water-cock. When the water is boiled in the large flask, the 
steam passes into the inner tube of the condenser. As this is 
surrounded by cold water, the steam condenses and the distilled 
water collects in the receiver. 



CHAPTER VI. 

CONSTITUTION OF MATTER. — ATOMIC THEORY.— 
ATOMS AND MOLECULES. — CONSTITUTION. — VA- 
LENCE. 

The Atomic Theory as proposed by Dalton, — We have 
seen how Dalton, at the beginning of this century, dis- 
covered the law of multiple proportions. This law, as well 
as that of definite proportions, called for explanation. The 
questions to be answered are: (1) Why do the elements 
combine in definite proportions ? (2) Why, when ele- 
ments combine with each other in more than one way, do 
the relative quantities which enter into combination in the 
different cases bear simple relations to one another ? and 
(3) What is the significance of the figures representing the 
combining weights ? Dalton saw that the facts referred 
to could be explained by the atomic theory, while, on the 
theory that matter is indefinitely divisible, they appear to 
be inexplicable. It is only necessary to assume that each 
element is made up of particles which are not divisible in 
chemical processes, and that these particles, or atoms, 
have definite weights. The atoms of any one element must 
be supposed to have the same weight, while the atoms of 
different elements have different weights. Now, when 
chemical combination takes place, Dalton supposed that 
the action is between the atoms. The simplest case is that 
in which combination takes place in such way that each 
atom of one element combines with one atom of another; 
but, besides this kind of combination, we may have that 
in which one atom of one element combines with two 
atoms of another, or two of one may combine with three 
of another, etc. Suppose two elements A and B, the 

97 






9^ COLLEGE CHEMISTRY. 

weights of whose atoms are to each other as 1:10, are 
brought together, and they combine in the simplest way, 
i.e., one atom of one with one atom of the other, then it 
is plain that in the compound AB the elements will be 
contained in the proportion of 1 part of A to 10 parts of 
B, whether a small or a large quantity of the compound 
is formed, and no matter in what proportions the elements 
are brought together. If they should be brought together 
in the proportion of their atomic weights (1 : 10), then no 
part of either element will be left uncombined after the 
act of combination has taken place. If, however, a larger 
proportion of either element is taken than that stated, 
then the quantity of the one which is in excess of this pro- 
portion will be left uncombined. This is in accordance 
with what we know takes place, and it is a conclusion 
drawn from the theory. No matter how many atoms of 
A we may take, the same number of atoms of B will be 
required to combine with all of them. But each atom of 
B weighs 10 times as much as each atom of A, therefore 
the total mass of B which enters into combination must 
be 10 times that of A with which it combines. It may 
be, however, that these same elements can form other 
compounds with each other. If A and B represent the 
atoms of the elements and we assume these atoms to be 
chemically indivisible, then the other compounds must be 
represented by such symbols as AB 2 , A 2 B, AB 5 , A 3 B, 
A^B^, etc., which represent compounds in which 1 atom 
of A is combined with 2 atoms of B\ 2 of A with 1 of B; 
1 of A with 3 of B; 3 of A with 1 of B; 2 of A with 3 
of B ; etc. ; or they also represent compounds in which 
1 part by weight of A is combined with 20 parts by weight 
of B; 2 of A with 10 of B; 1 of A with 30 of B; 3 of A 
with 10 of B; 2 of A with 30 of B; etc. It is therefore 
clear that, if the atomic theory as put forward by Dalton 
is true, the elements must combine according to the laws 
of definite and multiple proportions; and it appears that 
the figures which represent the combining weights of the 
elements must either bear to one another the same relation 



USE AND VALUE OF A THEORY. 99 

as the weights of the atoms, or the atomic weights or the 
relative weights of the atoms must be closely related to 
the combining weights, as will be pointed out more clearly 
presently. 

Use and Value of a Theory. — The relation of a theory 
to facts is very simple, but is frequently misunderstood. 
The relation may be conveniently illustrated by the case 
under consideration. By a careful investigation of a 
number of chemical compounds it was shown that in each 
of them the same elements always occur in the same 
proportion. This led to the belief that this is true of 
every chemical compound, and after further investigation 
which, as far as it went, show r ed the surmise to be correct, 
the law of definite proportions w T as proposed. This law is 
simply a statement of what has been found true in all cases 
examined. It involves no speculation. It is a statement 
of fact. It may be said that the statement or law must 
be open to some doubt for the reason that all possible cases 
have not been examined, and it may not hold true for some 
of these unexamined cases. The reply to this is that it has 
been found true in a very large number of cases and in all 
cases which have been investigated. It is true for the 
present state of our knowledge, and that is all we can 
demand of any law. Again, further investigation led to 
the discovery of the law of multiple proportions, which is 
also a statement of what has been found true in all cases 
investigated. It, like the law of definite proportions and 
in the same sense, is a statement of fact. But having gone 
thus far, we now ask, what is the explanation of these 
laws ? We know the facts — what is the explanation ? By 
experiment we cannot go beyond these facts, but it is 
possible to imagine a cause and then proceed to see 
whether the imagined cause is sufficient to account for the 
facts. This is what Dalton did. He imagined that 
matter is made up of atoms of definite weights, and that 
chemical combination takes place in simple ways between 
these atoms. This imagined cause is the atomic theory. 
It is not a statement of anything found by investigation. 
L. of C. 



ioo COLLEGE CHEMISTRY. 

It is not a statement of an established fact. It may or 
may not be literally true, but at all events it is the best 
guess that has ever been made as to the cause of the 
fundamental laws of chemical action, and it furnishes a 
very convenient means of interpreting the facts of chemis- 
try. Since the atomic theory was first proposed it has been 
accepted by nearly all chemists. It has been of great value 
in suggesting methods of work, and has contributed largely 
to the advance of chemistry. Any theory which is in 
accordance with the facts and leads to the discovery of new 
facts is of value, whether it should eventually prove to be 
true or false. At the same time a false theory may do 
much harm, as it may lead men to misinterpret the facts 
which they observe, and thus retard progress. 

Atomic Weights and Combining Weights. — If the 
atomic theory is true, the atoms of each element must have 
definite weights, and the determination of these atomic 
weights must evidently be of great importance. By analy- 
sis of compounds we can only determine the proportions 
by weight in which the elements combine with one another. 
Can we in this way determine the atomic weights ? In the 
first place, it is clear that it is out of the question to think 
of determining the absolute weights of the atoms, and all 
that we can possibly do is to determine their relative 
weights. As of all the elements hydrogen enters into 
combination in smallest relative quantity, its atomic weight 
is taken as the unit of the system, and the problem before 
us is to determine how many times heavier the atoms of 
the other elements are than that of hydrogen. If every 
element combined with hydrogen in only one proportion 
the problem would be a comparatively simple one. Thus 
the three elements chlorine, bromine, and iodine combine 
with hydrogen, forming only one compound each. On 
analysis these are found to contain respectively 

1 part of hydrogen to 35.18 parts of chlorine; 

1 " " " 79.36 " " bromine; and 

1 " " " 125.90 " " iodine. 



ATOMIC WEIGHTS AND COMBINING WEIGHTS. 101 

There is no reason for believing that in these compounds 
the elements are combined in any but the simplest way, 
i.e., that each atom of hydrogen is combined with one 
atom of chlorine to form hydrochloric acid, etc. If this 
is true, then the atom of chlorine must weigh 35.18 times, 
that of bromine 79.36 times, and that of iodine 125.90 
times as much as that of hydrogen, or, in other words, the 
atomic weights of chlorine, bromine, and iodine are 
respectively 35.18, 79.36, and 125.90. It will, however, 
be observed that there is no evidence as to whether the 
elements in these compounds are combined in the simplest 
way or not. It is possible, as far as we know, that one 
atom of hydrogen may combine with two or three of chlo- 
rine, or that one of chlorine may combine with two or 
three of hydrogen. As there is, however, no evidence 
upon this point the simplest assumption is made. 

If we take the case of oxygen the problem is more com- 
plex. In water the elements are combined in the propor- 
tion of 7.95 parts by weight of oxygen to 1 of hydrogen, 
and from this we should naturally conclude that the atomic 
weight of oxygen is 7.95; but further study shows that 
this conclusion is not justified. Hydrogen and oxygen 
form a second compound known as hydrogen dioxide or 
hydrogen peroxide in which there are 15.90 parts by 
weight of oxygen to 1 of hydrogen. This may be explained 
in the terms of the atomic theory by assuming that water 
is represented by the formula HO, and hydrogen dioxide 
by H0 2 . But it may be that the atomic weight of oxygen 
is 15.89, and then water must be represented by the 
formula H 2 0, and hydrogen dioxide by HO. Simple 
analysis of the compounds is not sufficient to enable us to 
decide between these possibilities. It is, therefore, evident 
that in order to determine the atomic weights something 
besides the determination of the composition of compounds 
is necessary. The figures representing the combining 
weights found in this way will, however, either be identi- 
cal with the atomic weights or will bear a simple numerical 
relation to them. 



102 COLLEGE CHEMISTRY. 

Molecules. — Investigation of certain phenomena of 
light, of electricity, of liquid films and the conduct of 
gases has led physicists to the conclusion that matter is 
not continuous, but made up of small particles, which are 
called molecules. A gaseous molecule is defined as " that 
minute portion of a substance which moves about as a whole, 
so that its parts, if it has any, do not part company during 
the motion of agitation of the gas." It would be out of 
place here to present the physical facts upon which the 
molecular theory rests. Suffice it to say that it is the only 
theory which has been found adequate to account for the 
behavior of gases. 

Avogadro's Hypothesis. — The fact that gases conduct 
themselves in the same way under the influence of changes 
in temperature and pressure can only be explained by 
assuming that equal volumes of all gases and vapors contain 
the same number of ultimate particles or molecules at the 
same temperature and pressure. 

This is a deduction from the well-tested molecular 
theory of gases. It was, however, originally put forward 
by the Italian chemist Avogadro from a study of chemical 
as well as of physical facts, and a little later it was sug- 
gested as probable by the French physicist Ampere. It is 
therefore generally spoken of as Avogadro's hypothesis, 
and sometimes as Ampere's hypothesis. Absolute proof 
of its truth cannot be given, but it is in thorough accord- 
ance with a large number of well-known facts, and it is 
undoubtedly true, if the molecular theory of matter is true. 
It may therefore be considered as furnishing a solid foun- 
dation for further conclusions bearing upon the problem 
of the determination of the atomic weights. 

Distinction between Molecules and Atoms. — If we con- 
sider any chemical compound, as water or hydrochloric 
acid, it is evident that the smallest particle or the molecule 
of the compound must be made up of still smaller particles. 
Thus, the smallest particle of water must contain smaller 
particles of hydrogen and oxygen, and the smallest particle 
of hydrochloric acid must contain smaller particles of 



MOLECULAR WEIGHTS. 103 

hydrogen and chlorine. These smallest particles of the 
molecules are the atoms. The molecules of the compounds 
are, according to this view, made up of the atoms of the 
elements. Similarly the elements themselves are, for good 
reasons which will be presented, believed to consist of 
molecules which are in turn made up of atoms of the same 
kind, though in a few cases the molecule of the element is 
identical with the atom. The difference between a com- 
pound and an element then is, in general, that the molecule 
of the compound consists of atoms of different kinds, while 
the molecule of an element consists of atoms of the same 
kind or, in a few cases, of one atom. Generally the atoms 
do not exict in the free or uncombined state, but, if they 
are set free by chemical action, they unite to form mole- 
cules. The following may serve as a definition of the con- 
ception of atoms at present held by chemists : 

Atoms are the indivisible constituents of molecules. They 
are the smallest particles of the elements that take part in 
chemical reactions, and are, for the greater part, incapable 
of existence in the free state, being generally found in com- 
bination with other atoms, either of the same hind or of 
different hinds. 

It cannot be too strongly emphasized that the views held 
in regard to the relations between molecules and atoms are 
based upon an enormous amount of painstaking study of 
facts, and in order fully to comprehend their value a study 
of most of these facts would be necessary. These views 
have gradually become firmly established as knowledge of 
the facts has grown more and more profound. Accepting 
them, we are now to see how they aid us in the problem 
with which we are -dealing, viz., the determination of the 
atomic weights. 

Molecular Weights. — If equal volumes of gases contain 
the same number of molecules at the same temperature 
and pressure, it is only necessary to determine the weights 
of equal volumes of gases to learn the relative weights of 
their molecules. Thus, if ~,?e weigh a litre of each of three 
gases, and find that the weights are to one another as 1 to 



104 COLLEGE CHEMISTRY. 

2 to 3, then it follows that the relation between the 
weights of the molecules of these gases is expressed by 
these figures, or, in other words, the molecule of the second 
gas is twice as heavy, and that of the third gas is three 
times as heavy as that of the lightest. The determination 
of the relative weights of the molecules of substances which 
either are gaseous or can be converted into gases resolves 
itself simply into a determination of the weights of equal 
volumes. In representing the molecular weights we may 
use any figures that are most convenient, provided only 
that they bear to one another the relations determined by 
experiment. If, however, we call the atomic weight of 
oxygen 16, then our system of molecular weights must be 
based upon this standard. But 16 is not the molecular 
weight of oxygen. In other words, the atom and the 
molecule of oxygen are not identical. This conclusion is 
drawn from a large number of facts. It is the result of a 
half century's work by the leading chemists of the world. 
One way in which it can be made clear is this: The weight 
of a given volume of oxygen gas as compared with the 
weight of the same volume of water in the form of gas 
under the same conditions of temperature and pressure is 
as 32 to 18.02. These equal volumes of gases contain the 
same number of molecules according to the hypothesis of 
Avogadro. Each molecule of water is made up of 1 atom 
of oxygen with the atomic weight 16 and 2 atoms of 
hydrogen with the atomic weight 1.01. If 16 is the atomic 
weight of oxygen, water cannot have a smaller molecular 
weight than 18.02. But, if the molecular weight of water 
is 18.02, that of oxygen must be 32, or it must be twice 
as great as the atomic weight. No matter what atomic 
weight we may adopt for oxygen this reasoning holds good. 
The standard molecule is therefore that of oxygen or, 
better, the basis of the system is a hypothetical gas having 
3*2 the Aveight of oxygen. This may be called the normal 
gas. Having adopted this standard, the molecular weight 
of a gas may be defined as its weight compared with that 
of th© normal gas. Below are given the molecular weights 



DEDUCTION OF ATOMIC WEIGHTS. 



105 



of a few elements and compounds which have been deter- 
mined by the method just described: 



Name. 


Molecular 
Weight. 


Molecular 
Formula. 


Hydrogen 


2.02 
28.08 
18.02 
36.46 
17.07 
16.04 
28 
44 


H 2 


N itrogen 


N a 


Water 


H 2 
HC1 
NH 3 


Hydrochloric acid 

Ammonia 


Marsh-gas 


CH 4 


Carbon monoxide 


CO 


Carbon dioxide 


C0 2 







Deduction of Atomic Weights from Molecular Weights. 

— The determination of molecular weights does not neces- 
sarily carry with it the determination of the atomic 
weights. It is plain from what has already been said that 
a knowledge of the molecular weight of an element does 
not convey a knowledge of its atomic weight. If, for 
example, w T e learn that the molecular weight of nitrogen 
is approximately 28, we have no means of judging from 
this what the atomic weight is. It is plainly necessary to 
know of how many atoms each molecule of nitrogen is 
made up, and to learn this is not a simple matter. It is 
easier to determine the atomic weight of an element 
through a study of its compounds. Suppose it is desired 
to determine the atomic weight of oxygen. We first 
determine the molecular weights of a number of com- 
pounds that contain oxygen, and then analyze these com- 
pounds. We then see what the smallest figure is that is 
required to express the weight of the oxygen that enters 
into the composition of the molecules, and this figure is 
selected as the atomic weight. The molecular weights and 
the composition of several oxygen compounds are given in 
the following table. The figures in the third column are of 
course determined by analysis, an example of the methods 
used having been given in the chapter on water. Stated in 
ordinary language, the figures in the case of carbon mon- 



io6 



COLLEGE CHEMISTRY. 





Com 


position. 


2.02 


parts hydrogen, 


16 


" 


oxygen. 


12 


" 


carbon, 


16 


a 


oxygen. 


12 


" 


carbon, 


32 


tt 


oxygen. 


14.04 


" 


nitrogen, 


16 


" 


oxygen. 


28.08 


' : 


nitrogen, 


16 


* ' 


oxygen. 


32.06 


>« 


sulphur, 


32 


" 


oxygen. 


32.06 


" 


sulphur, 


48 


t< 


oxygen. 



Compound. Mol. Wt. Approx. 
Water , 18.02 

Carbon monoxide 28 

Carbon dioxide 44 

Nitric oxide 30.04 

Nitrous oxide 44.08 

Sulphur dioxide 64.06 

Sulphur trioxide 80.06 



oxide mean that the molecule of this compound weighs 28 
times as much as the normal gas, and the 28 parts by weight 
of matter are made up of 12 parts of carbon and 16 parts 
of oxygen. Considering now the composition of the com- 
pounds in the table, it will be seen that the smallest mass 
of oxygen that enters into the composition of any of the 
molecules weighs 16 in terms of the normal gas. We find 
twice this mass, as in carbon dioxide and sulphur dioxide; 
and three times, as in sulphur trioxide, but no smaller 
mass. Now, an examination of all compounds of oxygen 
that are known to us in the form of gas or vapor shows 
the same thing to hold true; that is to say, the smallest 
mass of oxygen that enters into the composition of mole- 
cules is 16 times as great as that of the normal gas. The 
conclusion is therefore drawn that 16 is the atomic weight of 
oxygen. The possibility that the atomic weight of oxygen 
is less than this figure is not excluded. It may be that in 
the simplest oxygen compounds now known there are two 
or more atoms of this element in the molecules. But in 
the total absence of evidence on this point all we can do 
is to accept the figure 16 as in perfect accordance with all 
our knowledge of oxygen compounds. 

In this way the atomic weights of all elements that form 
gaseous compounds or compounds that can be converted 



MOLECULAR FORMULAS. 107 

into vapor have been determined; and the determinations 
made in this way are regarded as the most reliable. 

Exact Atomic Weights determined by the Aid of 
Analysis. — By determining molecular weights it is possible 
to decide approximately what figure represents the atomic 
weight of an element, but the methods employed in mak- 
ing determinations of molecular weights are liable to slight 
errors, and therefore the atomic weights obtained directly 
from the molecular weights deviate slightly from the true 
figures. In order to determine the atomic weights with 
the greatest possible accuracy, the most refined methods 
of chemical analysis are brought into play, and the figures 
in the table on page 19 have been determined in this way 
by a combination of a study of the specific gravity of gases 
and by the most careful analyses, together with some other 
methods which will be taken up later. 

Molecular Formulas. — The symbols of chemical com- 
pounds first used were intended to express simply the 
composition of the compounds, and this can be done as 
was explained in Chapter I, by adopting a system of com- 
bining weights of the elements. According to the theory 
explained in the last chapter, the smallest particle of every 
compound is a molecule, and each molecule is made up of 
atoms. It appears, therefore, desirable for the sake of 
uniformity that the symbols used to represent chemical 
compounds should represent molecules. Where the molec- 
ular weight of a compound, the atomic weights of the 
elements of which it is composed, and its composition are 
known, there is no difficulty in representing it by a molec- 
ular formula. Thus, the molecular weight of ammonia 
is found by experiment to be approximately 17, and the 
17 parts are made up of 11 parts of nitrogen and 3 parts 
of hydrogen. The atomic weight of nitrogen is found by 
the method which has just been described to be very nearly 
11. Therefore the molecule of ammonia weighing 17 
parts is composed of 1 atom of nitrogen weighing 14 parts 
and 3 atoms of hydrogen weighing 3 parts. The compo- 
sition of the molecule is therefore represented by the 



108 COLLEGE CHEMISTRY 

formula NH 3 . Similarly the composition of the molecule 
of water is represented by the formula H 2 0; that of hydro- 
chloric acid by HOI; that of marsh-gas by CH 4 ; etc., etc. 
Every formula now in use is intended to represent a mole- 
cule of the compound for which it stands. In regard to 
the molecular weights of compounds that are not gaseous 
nor convertible into vapor, Avogadro's method is plainly 
of no avail. Methods have, however, been devised which 
are applicable to a number of these (see Chapter XXIII). 
Constitution. — When hydrochloric acid is formed, we 
conceive that each atom of hydrogen combines with one 
atom of chlorine, and that the molecules of the resulting 
compound are made up each of an atom of hydrogen and 
an atom of chlorine. What the act of combination con- 
sists in we do not know. We simply know that something 
very remarkable takes place, and that as a consequence the 
hydrogen and chlorine cease to exist in their original 
forms. It is idle at present even to speculate in regard to 
the character of the change. The fact of union is ex- 
pressed by writing the symbols of the elements side by side 
without any sign between them, as HC1, or, sometimes, it 
is convenient to use a line to indicate chemical union, 
thus: H-Cl. According to the molecular theory the 
molecule of water consists of two atoms of hydrogen and 
one of oxygen, as represented by the formula H 2 0, and the 
question now suggests itself whether all three atoms are in 
combination with one another or whether each of the 
hydrogen atoms is in combination with the oxygen atom, 
but not with each other, as represented by the formula 
H-O-H. So, too, in the case of ammonia, the molecular 
formula of which is NH 3 , the question suggests itself : Are 
the three atoms of hydrogen in combination with the atom 
of nitrogen, but not with one another, as represented in 

the formula N^-H ? It is extremely difficult to answer 

such questions, but, nevertheless, certain facts are known 
which enable us to draw probable conclusions. Formulas 



VALENCE. 109 

which express the composition of molecules and at the 
same time express the relations or the connections which 
exist between the atoms are called constitutional formulas. 
These constitutional formulas are very frequently used at 
present, but sometimes without a sufficient basis of facts 
to justify them. Whenever they are used in this book, 
the reasons for them will be stated as fully as may appear 
necessary. 

Valence. — The formulas of the hydrogen compounds of 
chlorine, oxygen, nitrogen, and carbon, all determined by 
the same method, are 

C1H OH 2 NH 3 CH 4 . 

A consideration of these formulas and of many similar 
ones has led to the belief that the atoms of different ele- 
ments differ in their power of holding other atoms in 
combination. The simplest explanation of the composition 
of the compounds above represented is that the atoms of 
chlorine, oxygen, nitrogen, and carbon differ in their 
power of holding hydrogen atoms in combination. Hy- 
drogen and chlorine combine in only one way, 1 atom of 
chlorine combining with 1 of hydrogen; 1 of oxygen com- 
bines with 2 of hydrogen; 1 of nitrogen with 3 of 
hydrogen; and 1 of carbon with 4 of hydrogen. The 
limit of the combining power of the atom of chlorine is 
reached when it has combined' with one atom of hydrogen. 
And as one chlorine atom can hold but one atom of 
hydrogen in combination, so one atom of hydrogen can 
hold but one atom of chlorine. Either the hydrogen atom 
or the chlorine atom may be taken as an example of the 
simplest kind of atom. Any element like hydrogen or 
chlorine is called a univalent element ; an element like 
oxygen whose atom can hold two unit atoms in combina- 
tion is called a bivalent element ; an element like nitrogen 
whose atom can hold three unit atoms in combination is 
called a trivalent element ; and an element like carbon 
whose atom can hold four unit atoms in combination is 
called a quadrivalent element. Most elements belong to 



no COLLEGE CHEMISTRY. 

one or the other of these four classes, though there are 
some which can hold five, six, and even seven unit atoms 
in combination. These are called quinquivalent , sexiva- 
le?it, and septivalent respectively. 

Valence is defined as that property of an element by 
virtue of which its atom can hold a definite number of 
other atoms in combination. In the formation of com- 
pounds the valence of the elements determines how many 
atoms of any element can enter into combination with any 
other. The atoms are sometimes spoken of as having 
bonds which are graphically represented by lines. Thus, 
a univalent element is said to have one bond, as repre- 
sented by H-, C1-, etc. ; a bivalent element is said to have 

two bonds, -0-, -S-, etc. ; a trivalent element three, -N-; 

and a quadrivalent element four, -0— Of course, this is 

i 
merely a symbolical representation of the idea that each 

atom has a definite power of combining with others. It 

is further said that when the atoms unite these bonds 

become satisfied. Thus when one atom of hydrogen unites 

with one of chlorine, the bond of each is regarded as 

uniting with the bond of the other, and this is represented 

by the symbol H-Cl. So, too, when two atoms of hydrogen 

unite with one of oxygen, the compound is represented in 

TT 

this way: H-O-H or 0<tt. In the union of atoms, 

further, to use the figurative language, two bonds may be 
satisfied by two univalent atoms or by one bivalent atom. 
Thus, in marsh-gas, CH 4 , the four affinities or bonds of 
the quadrivalent carbon atom are regarded as being satis- 
fied by the four univalent hydrogen atoms. In the com- 

H 

i 
pound H-C — 0, however, two of the bonds are regarded 

as satisfied by two univalent hydrogen atoms, and the 

other two by one bivalent oxygen atom; and, again, in the 

compound = C = 0, the four bonds of the carbon atom 

are regarded as satisfied by two bivalent oxygen atoms. 






REPLACING POWER OF ELEMENTS III 

Replacing Power of Elements. — As has been seen, when 
potassium acts upon water the action is represented by the 
equation 

K + H 2 = KOH + H. 

Expressing this reaction by means of formulas which take 
the valence of the elements into consideration we have the 
following- : 

K -f H-O-H = K-O-H + H. 



l o 



According to this, one atom of potassium is substituted 
for one atom of hydrogen in water, and in the compound 
formed, the atom of potassium is regarded as occupying 
the place of the hydrogen atom. The elements calcium 
and barium are bivalent like oxygen, as shown in the 
compounds CaCl 2 and BaCl 2 , in which the atom of calcium 
as well as that of barium holds two univalent atoms of 
chlorine in combination. When these elements act upon 
water, hydrogen is liberated as with potassium, but each 
atom replaces two atoms of hydrogen, as represented in the 
equations 

Ca + 2H 2 = CaH 2 2 4-2H; 
Ba + 2H 2 = BaH 2 2 + 2H. 

Or expressing the changes by valence formulas, we have 

r , H-O-H p O-H . w 
Ca + H-O-H = Ca< _ H + H 2 ; 

B , H-O-H ^ O-H . jr 
Ba + H-O-H = Ba < O-H + H *' 

A trivalent element. acting in the same wav would give a 

•O-H 
compound of the general formula M^— O-H, in which M 

\0-H 
represents any trivalent element. 

80, too, in the action of various elements upon hydro- 
caloric acid, HOI, a univalent element like sodium or 
potassium replaces one hydrogen atom in one molecule of 



H2 COLLEGE CHEMISTRY. 

the acid, and forms a compound of the general formula 
MCI, as, for example, KC1 and NaCl. A bivalent ele- 
ment, like zinc, replaces two atoms of hydrogen in two 
molecules of the acid, forming the compound ZnCl 2 : 

Zn + HCl =Zn< Ci+ H2 ' 

A trivalent element forms a compound of the general 
formula MC1 3 , as, for example, aluminium chloride : 

/01 

Alf-Cl. 

xa 

The above explanation of the hypothesis of valence will 
suffice by way of introduction to the subject. The relation 
which it bears to the facts is simply this: On studying 
the composition of chemical compounds we find that, in 
general, the elements combine with one another in com- 
paratively few proportions. Thus, hydrogen and chlorine 
combine with each other in only one proportion ; hydrogen 
and oxygen in two; nitrogen and hydrogen in four; etc. 
There is something limiting the complexity of compounds. 
We might study the laws governing the complexity of 
compounds without any hypothesis as to the cause, but 
the hypothesis of valence is a convenient explanation of 
these laws, and it has been of much service in furnishing 
chemists with a simple language for representing chemical 
changes. 

EXPERIMENTS. 

It would be well in this connection to determine the specific 
gravity of some substance in the form of vapor. The principal 
methods for this purpose are those of Dumas, Gay-Lussac, Hof- 
mann, and Victor Meyer. That of Dumas, which consists in 
measuring the volume and determining the weight of the vapor 
under observation, is the most accurate. The method of Hofmann 
is a modification of that of Gay-Lussac. It consists in weighing a 
small quantity of the liquid the specific gravity of whose vapor is 
to be determined, and, after introducing the liquid in a minute 
glass vessel into a eudiometer over mercury, heating the eudiom- 



METHOD OF DUMAS. 



1 I~s 



cter and its contents by passing steam through a jacket surround- 
ing it and measuring the volume of vapor formed. The method 
of Victor Meyer is used when it is required to determine the spe- 
cific gravity of the vapor of a substance which boils at a high 
temperature. 

Method of Dumas. 

Experiment 61. — In this method the liquid to be vaporized is 
brought into a small balloon like that shown in Fig. 29. The dry 
balloon is first weighed, and a small quantity of liquid then intro- 
duced by gently heating the balloon and putting the point of its 
stem into the liquid, when, on cooling, the liquid rises and enough 
is easily brought into the balloon in this way. The balloon is now 
placed (in the position shown in Fig. 30) in a bath of water, oil, 





Fig. 29. 



Fig. 30. 



or paraffin, according to the boiling-point of the liquid. The bath 
is heated 30-40° above the boiling-point of the liquid under ex- 
amination. The air is thus driven out and the balloon is filled 
with the vapor. When vapor no longer escapes, the point of the 
stem is closed by melting it with a mouth blowpipe. The balloon 
is then cleaned, dried, and weighed. The temperature of the bath 
and the height of the barometer are observed at the time the 
balloon is closed. The point of the stem is broken off under mer- 
cury, when the mercury rises and fills the balloon. By pouring 
the mercury out into a graduated cylinder the capacity of the 
balloon is determined. The specific gravity of the vapor is calcu- 
lated by the aid of the formula 

(#i — B + p) . (1 + 0.00366 . U) . 760 



D 



v . 0.001293 . hi 



H4 

in which 



COLLEGE CHEMISTRY. 



B = weight of balloon at t° and h mm.; 
Bx = " " " with vapor, at ti° and hi mm. 
v = capacity of the balloon in cubic centimetres ; 
0.001293 = weight of 1 cc. air at 0° and 760 mm.; 
p = weight of air in balloon at t° and h mm. 







Method of Victoe Meyer. 

Experiment 62. — In this method a known weight of substance 
is converted into vapor, and the volume of vapor formed is deter- 
mined by measuring the volume of air which it 
displaces. The apparatus consists of an outer cyl- 
indrical vessel A, Fig. 31, and an inner vessel B, 
which is connected with a tube C. The vessel B 
has a capacity of about 100 cc, and is about 200 
mm. long. The tube (7, with its funnel-shaped 
end JE, is about 600 mm. long. First, a small 
quantity of some substance with a boiling-point 
high enough to secure the complete conversion 
into vapor of the substance to be studied, is put in 
the bottom of the vessel A, and a little ignited 
asbestos or dry mercury in the bottom of the ves- 
sel B. The substance in A is now heated to boiling, 
and E is closed with a rubber stopper. After a 
time the temperature of the air in B is raised to 
that of the vapor in A, and no more escapes from 
the tube D. When this condition of equilibrium 
is reached, a small weighed quantity of the sub- 
stance under examination is dropped into the 
vessel B, the stopper being removed from E and 
quickly replaced. The substance is converted into 
vapor, and displaces an equivalent volume of air, 

Qand this displaced air is collected over water in 
the measuring-tube placed over the end of D. 
Fig. 31. When no more air escapes, the volume is deter- 

mined in the usual way. The specific gravity of the substance is 
calculated by the aid of the following formula : 



B = S x 



(1 + 0.00366 x t) 760 



(B- w) V x 0.001293' 

in which 0.001293 is the weight of 1 cc. air in grams at 760 mm. 
and 0° ; and, further, 



METHOD OF VICTOR MEYER. 115 

.S v = weight of substance taken ; 
t = temperature of the room, or of the water in the measuring 

apparatus ; 
]> = height of barometer ; 
u: = tension of aqueous vapor ; 
V — observed volume of air ; 

or, the formula can be simplified by division, when it takes this 
form : 

(1 + 0.00366 x £)587,780 



D=S 



(B —w) V 



The above is the simplest form of apparatus used. To avoid 
opening and shutting the vessel in order to introduce the sub- 
stance, an arrangement has been devised for holding the sub- 
stance below the stopper until the proper temperature is reached, 
and then releasing it without disturbing the stopper. 



CHAPTER VII. 

OZONE.— ALLOTROPY.— NASCENT STATE.— HYDROGEN 
DIOXIDE. 

Occurrence. — Ozone has long been thought to be present 
in the air in small quantity, but careful research has made 
this occurrence appear doubtful. 

Preparation.— It was observed iu the last century that, 
when a powerful electric machine is worked iu a room, 
something is formed which has a strong odor, and the same 
odor was noticed during thunder-showers. Afterwards the 
same thing was noticed when water is decomposed by an 
electric current, and when phosphorus is exposed to moist 
air. The substance was at first supposed to be a compound 
of water and oxygen, but long- continued investigation 
showed that it could be made from pure dry oxygen, and 
that by heat and other means it is converted into oxygen. 

It is prepared mixed with oxygen by passing electric 
sparks through ordinary oxygen in an apparatus con- 
structed on the principle of that represented in Fig. 32. 




Fig. 32. 

AA is a glass tube about an inch in diameter closed at 
the ends with brass caps or with corks, covered with 
shellac on the inner side. A metallic cylinder BB covered 
with tin-foil is placed inside this tube. The cylinder is 

V 116 



OZONE—PROPER TIES. 1 1 7 

connected with the tubes CO, and through these and the 
cylinder a current of cold water is kept flowing. The 
oxygen passes through the glass tube by means of the tubes 
DjD, and necessarily passes through the narrow space 
between the glass tube and the metallic cylinder. Around 
the outside of the glass tube is wound a strip of tin-foil. 
By means of the wires F and E connection is established 
with the poles of an induction-coil, or a Holtz electrical 
machine. 

Ozone is also made by placing a few pieces of phosphorus 
in the bottom of a good-sized bottle and partly covering 
them with water; and, finally, it is made by treating 
barium dioxide, Ba0 2 , and some other compounds rich in 
oxygen with sulphuric acid. 

Properties. — Ozone is a gas which can be condensed to 
the liquid form, the liquid having a blackish-blue color. 
It has a strong odor, and acts in an irritating way upon 
the membranes lining the throat. Its chemical conduct 
is entirely different from that of oxygen. While the latter 
at ordinary temperatures is not an active substance, ozone 
is. It acts upon most substances of animal or vegetable 
origin, and oxidizes nearly all the metals, besides produc- 
ing a variety of other changes which are not produced by 
oxygen at the ordinary temperature. Among the charac- 
teristic changes which may be made use of for the purpose 
of detecting ozone are the following : (1) It liberates iodine 
when brought in contact with potassium iodide ; the action 
being represented thus: 

2KI + H 2 + = 2KOH + 21. 

Now, iodine has the power to turn starch blue. There- 
fore, if ozone is brought into a solution containing starch 
and potassium iodide, a blue color is produced. (2) Ozone 
combines with metallic silver to form silver peroxide, 
which is brown. In order to detect ozone, therefore, a 
strip of polished silver may be exposed to the gas, and if 
it turns brown ozone is present. 



n8 



COLLEGE CHEMISTRY. 



fk 



When heated to 300° ozone loses its characteristic odor 
and is converted into ordinary oxygen. It thus appears 
that we can start with ordinary oxygen, and by means of 
an electric current convert it into a substance with much 
more active chemical properties and differing from it so 
markedly that one would hardly suspect the close relation 
between the two; and then, further, by simply passing the 
active substance, or ozone, through a tube heated to 300°, 
it is converted into oxygen without loss of weight. 

Relation between Oxygen and Ozone. — When experi- 
ment had shown that oxygen and ozone are convertible 
one into the other without change of weight, the sugges- 
tion was made that the difference between them might be 
due to a difference in the number of atoms 
of oxygen contained in the molecule of each. 
It might be, for example, that in the mole- 
cule of oxygen there are two atoms and in 
that of ozone three or more. If this view is 
correct there should be a difference between 
the specific gravities of the two gases, and 
by a study of this difference it should be 
possible to draw a conclusion as to the con- 
stitution of the molecule of ozone. That 
there is a change of volume when oxygen is 
changed to ozone was shown by enclosing the 
former in a tube constructed as shown in 
Fig. 33. The large part of the tube is fur- 
nished with two platinum wires which pass 
through the glass. By means of these wires 
a silent discharge of electricity is kept up 
through the oxygen, and it is thus partly converted into 
ozone. In the smaller bent part of the tube there is a 
small column of concentrated sulphuric acid which serves 
as a stopper. Any change in the volume of the gas in the 
tube will cause this sulphuric acid to change its position. 
Now, during the conversion of the oxygen into ozone it 
was noticed that the volume of the gas decreased, and that, 
on heating the tube and thus converting the ozone into 



/ 



4^i 



Fig. 33. 



OZONE IN THE AIR. 119 

oxygen again, the original volume of the latter was 
restored. Unfortunately for this purpose it is not possible 
by any means to convert more than a comparatively small 
proportion of the oxygen into ozone, and it is not possible 
in the experiment described to determine the relation 
between the decrease in volume and the extent of the con- 
version. 

The most satisfactory determination of the relation 
between the weights of equal volumes of oxygen and of 
ozone has been made as follows: A mixture of ozone and 
oxygen containing about 8 to 9 per cent of ozone was 
liquefied by placing the vessel in which it was contained in 
liquid air. After the vessel was filled with the mixture it 
was raised above the liquid air so that its temperature 
became somewhat higher. Under these circumstances the 
oxygen boiled off, leaving a little ozone. More of the mix- 
ture was then liquefied in the same vessel, and the oxygen 
again allowed to boil off. In this way a small amount of 
a blackish-blue opaque liquid was obtained, which was 
probably the purest specimen of ozone that has ever been 
made. This liquid was then converted into gas, and the 
weight, as compared with that of the same volume of 
oxygen, determined. The ratio was found to be 1.3698 : 1. 
But this gas still contained oxygen, and an analysis showed 
that only 86. 16 per cent of the mixture was ozone. Making 
allowance for this, it appears that the weight of a given 
volume of oxygen to that of the same volume of ozone is 
1:1.5. Interpreting this fact in terms of the molecular 
theory, it is clear that, if the molecular weight of oxygen 
is 32, that of ozone is 48. Or, further, if there are two 
atoms in the molecule of oxygen, there are three in that 
of ozone. These conclusions are represented by the sym- 
bol 2 for oxygen and 3 for ozone. 

Ozone in the Air. — No satisfactory evidence of the 
occurrence of ozone in the air has ever been furnished. 
Very little, if anything, is known in regard to the effect of 
ozone on the health. Larger quantities would undoubtedly 
act injuriously. It is commonly believed that small quan- 



120 COLLEGE CHEMISTRY. 

tities are advantageous, as it tends to destroy substances 
which are unwholesome. This subject has not, however, 
been studied with sufficient care to justify a positive 
opinion in regard to it. The difficulty in drawing a con- 
clusion is increased by the fact that there are other sub- 
stances present in the air which resemble ozone in some 
respects, as, for example, hydrogen dioxide, H 2 2 (which 
see). 

Allotropy. — The occurrence of an element in two or 
more different modifications is called allotropy. Thus, 
ozone is called an allotropic form of oxygen. There are, 
however, a number of other cases of a similar kind. 
Phosphorus, for example, presents itself in three or four 
different varieties which, in some respects, differ markedly 
from one another. Carbon also appears in three different 
forms. Whether in all cases the difference between the 
allotropic forms of an element is due to a difference in the 
molecular constitution, as in the case of oxygen, is im- 
possible to decide at present, for the reason that the 
molecular weights of the substances cannot always be 
determined. The facts learned in regard to the relations 
betw r een oxygen and ozone make it appear quite probable 
that the explanation which holds good for this case will 
probably hold good for others. 

Varying Number of Atoms in the Molecules of one and 
the same Element. — It has been pointed out that the 
molecules of hydrogen, oxygen, and some other elements 
consist of two atoms each. We have just seen that the 
molecule of ozone consists of three atoms. There are some 
elements which contain a different number of atoms in 
their molecules, according to the temperature. Thus, 
sulphur is a solid substance which boils at 440° 0. The 
specific gravity of its vapor between 450° and 500° leads to 
the molecular weight about 192. But the atomic weight 
of sulphur is very nearly 32, therefore the molecule of sul- 
phur just above its boiling-point would appear to be made 
up of six atoms. As the temperature is raised, the specific 
gravity becomes less, and above 800° it corresponds to the 



NASCENT STATE. 12 1 

molecular weight 04, showing that at this high tempera- 
t are the molecule apparently consists of two atoms. Above 
that point the specific gravity does not change materially. 
According to the latest investigations on this subject, it 
appears probable that at low temperatures the molecular 
weight of sulphur in the form of vapor is 256, or, in other 
words, that there are eight atoms in the molecule. Facts 
somewhat similar to those just mentioned have been 
brought to light in the case of iodine, only here the vapor 
at the lower temperature has a specific gravity showing 
that the molecule consists of two atoms, but at very high 
temperatures the specific gravity shows that the molecule 
contains only one atom. The atomic weight of mercury 
is the same as its molecular weight, or, in other words, 
the molecule and atom are in this case identical. 

Nascent State. — One of the most curious phenomena 
connected with chemical action is that of the nascent state. 
Some elements, which under ordinary circumstances are 
inactive, show themselves possessed of marked activity if 
they are allowed to act the instant they are set free from 
their compounds. Thus, if carbon monoxide, CO, and 
ordinary oxygen are brought together at the ordinary tem- 
perature, no action takes place. If, however, carbon 
monoxide is brought in contact with something which 
easily gives off oxygen, it is converted into carbon dioxide. 
At ordinary temperatures, for example, chromium trioxide 
gives up oxygen to carbon monoxide as represented thus : 

3CO + 2Cr0 3 = 3C0 2 + Cr 2 3 . 

So, too, hydrogen at ordinary temperatures is an inactive 
element, but, if brought in contact with substances the 
instant it is set free from a compound, it produces many 
marked changes. Many substances w T hich would be left 
unchanged, if hydrogen were passed over them at ordinary 
temperatures, are changed if put in the liquid from which 
the hydrogen is being evolved. The most plausible ex- 
planation of facts like these is to be found in the molecular 
theory. Free oxygen and free hydrogen consist of mole- 



122 COLLEGE CHEMISTRY. 

cules made up of atoms which are in combination with 
each other. Before these molecules can act chemically 
they must be broken up into atoms. When carbon mon- 
oxide is brought together with oxygen in the molecular 
state, the condition is represented thus : 

CO + 2 . 

Before combination can take place the molecule, 2 , must 
be decomposed as represented thus : 

2 = O + O; 

and the atoms of oxygen then combine with the carbon 
monoxide. Now there are some substances which at 
ordinary temperatures yield oxygen atoms more readily 
than ordinary oxygen does. This is true of chromium tri- 
oxide, Cr0 3 , which in contact with substances that have 
the power to combine with oxygen breaks up thus : 

2Cr0 3 = Or 2 3 +0 + + 0. 

When hydrogen is liberated from a compound by chem- 
ical action the uncombined atoms are believed to be given 
off first; if there is nothing present with which the 
hydrogen can combine, these atoms combine with each 
other in pairs to form the molecules of ordinary hydrogen. 

The examples furnished by allotropy and the phenom- 
ena of the nascent state show how chemical facts which 
otherwise appear entirely incomprehensible are explained 
by the aid of the molecular theory. 

Hydrogen Dioxide or Hydrogen Peroxide, H 2 2 . — When 
barium dioxide is treated with hydrochloric acid or dilute 
sulphuric acid the reactions represented by the following 
equations take place: 

Ba0 2 + 2HC1 = BaCl 2 + H 2 2 ; and 
Ba0 2 + H 2 SO, = BaS0 4 + H 2 2 . 

In the latter case the compound BaS0 4 , or barium sul- 
phate, is insoluble, and the liquid can easily be separated 



HYDROGEN DIOXIDE. i?3 

from it by filtration. If the liquid is boiled, decomposi- 
tion of the hydrogen dioxide takes place, oxygen being 
liberated and water left behind. The decomposition is 
represented thus: 

2H 2 2 = 2H 2 + 2 ; r H 2 2 = H 2 + 0. 

The dioxide can be obtained in the form of a colorless 
liquid by allowing the solution to stand in a vacuum over 
sulphuric acid, or by distilling its solution in a vacuum. 
It boils at 8-i°-85° under a pressure of 68 mm. 

Properties. — Hydrogen dioxide is a clear, syrupy liquid 
which does not solidify at — 30°. It is characterized by 
marked instability. It breaks down slowly even at ordinary 
temperatures if simply allowed to stand. The decomposi- 
tion takes place easily under the influence of heat, and if 
heated rapidly to 100° explosion is apt to take place. The 
products are water and oxygen, as indicated above. In 
consequence of the ease with which hydrogen dioxide gives 
off oxygen it is a good oxidizing agent, and it is now 
manufactured on the large scale for use in bleaching and 
in medicine, and comes into the market in solutions of 
different strengths. The solution has a bitter taste; if 
concentrated, it affects the skin, causing a pricking sensa- 
tion, and making white spots, which, however, disappear 
in a few hours. 

Occurrence in the Air. — That hydrogen dioxide occurs 
in the air has already been stated. It is also found in rain 
and snow. The quantity in the air is extremely small, 
and it varies at different times of the day, the action of 
sunlight being evidently favorable to its formation. 

Characteristic Reactions. — Like ozone, hydrogen di- 
oxide decomposes potassium iodide, setting iodine free: 

H 2 2 + 2KI = 2KOH + I 2 . 

This fact may be utilized for the purpose of detecting 
the compound. The separation of the iodine does not 
take place readily as in the case of ozone, but the action is 



124 COLLEGE CHEMISTRY. 

hastened by the addition of a very small quantity of a 
dilute solution of ferrous sulphate, FeS0 4 . — An acid solu- 
tion of potassium permanganate is decolorized by hydrogen 
dioxide. — If in a glass cylinder a layer of ether is poured 
upon a solution of hydrogen dioxide and a drop of a solu- 
tion of potassium dichromate is then added, and the 
cylinder thoroughly shaken, the ether will take up a blue 
compound, and will itself become blue. 

When hydrogen dioxide is brought together with sub- 
stances which give up oxygen readily, action generally 
takes place involving decomposition of the hydrogen 
dioxide as well as of the other substance. Thus when it 
is brought together with silver oxide, Ag 2 0, this reaction 
takes place: 

A & + H 2 2 = Ag 2 + H 2 + 2 . 

So, also, it undergoes decomposition with ozone as repre- 
sented thus : 

3 + H 2 2 = H 2 + 20 2 . 

EXPERIMENTS. 

Ozone. 

Experiment 63. — Put a few sticks of ordinary phosphorus on 
the bottom of a good-sized bottle with a wide mouth, and partly 
cover the phosphorus with water. In a short time the odor of 
ozone will be perceptible, and the gas can also be detected by 
means of strips of paper which have been moistened with a dilute 
solution of potassium iodide and starch-paste. See whether such 
papers are changed in the air ? What is the cause of the change ? 
If convenient, examine the air in the neighborhood of a frictional 
electrical machine, and see whether it causes the papers to 
change color. 

Hydrogen Dioxide. 

Experiment 64. — Finely powder some barium dioxide, and add 
some of it to dilute sulphuric acid. Filter from the precipitated 
barium sulphate, and with the solution try the following reac- 
tions : 



HYDROGEN DIOXIDE: EXPERIMENTS. 125 

Heat some in a test-tube. What takes place ? Add to another 
small portion a little of a dilute solution of potassium permanga- 
nate. To another portion add a little finely powdered manganese 
dioxide. What is given off ? To a dilute solution contained in a 
small stoppered cylinder add a few drops of a dilute solution of 
potassium dichromate, and quickly add ether, and shake the 
cylinder thoroughly. 



CHAPTER VIII. 
CHLORINE.— HYDROCHLORIC ACID. 

Historical. — Sodium chloride or common salt, which is 
the principal compound of chlorine found in nature, has 
been known from the earliest times. In 1774 Scheele first 
called attention to chlorine in his treatise on the black 
oxide of manganese or manganese dioxide. In accordance 
with the ideas then prevailing, he called it ' ' dephlogisti- 
cated muriatic acid." Berthollet suggested in 1785 that 
it is oxidized hydrochloric acid, and it was then regarded 
as consisting of the hypothetical element, murium, in 
combination with oxygen. In 1810 Davy pointed out that 
the idea, previously expressed by Gay-Lussac and Thenard, 
that the substance is an element is in the highest degree 
probable, and he gave it the name chlorine (from jA oppds, 
greenish-yellow). Since that time everything learned in 
regard to chlorine has gone to show that it is an element. 

Occurrence of Chlorine. — Though widely distributed in 
nature, chlorine never occurs in the uncombined state, for 
the reason that it combines with other substances with 
great ease, and, if it were set free, it would at once enter 
into combination. It does not occur in very large quantity 
as compared with oxygen and hydrogen. It is found 
chiefly in combination with the element sodium as common 
salt, or sodium chloride, NaCl. It is also found in com- 
bination with other elements, as potassium, magnesium, 
etc., as in the celebrated mines at Stassfurt, Germany. 
In comparatively small quantity it occurs in combination 
with silver, forming one of the most valuable silver ores. 

126 



PREPARATION OF CHLORINE. 127 

All the chlorine we have to deal with is made from common 
salt. 

Preparation. — The problem to be solved in the prepara- 
tion of chlorine from common salt is the separation of the 
two elements sodium and chlorine. This cannot be 
accomplished directly as the separation of mercury and 
oxygen in the decomposition of mercuric oxide, HgO, and 
there is no easily obtained compound that gives off chlorine 
when heated. The method adopted consists in making 
hydrochloric acid, HC1, from sodium chloride, and then 
treating this acid with some substance which readily gives 
off oxygen. The change of sodium chloride to hydro- 
chloric acid is readily accomplished by treating salt with 
ordinary sulphuric acid — a reaction carried on on the large 
scale in the manufacture of sodium carbonate or "soda." 
When the two are brought together a change takes place 
which will be studied more in detail farther on. The 
reaction is represented by the equation 

(1) 2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1. 

Sodium Sulphuric Sodium Hydrochloric 

chloride acid. sulphate acid. 

As will be seen, the sodium of the sodium chloride and 
the hydrogen of the sulphuric acid exchange places. 

The decomposition of the hydrochloric acid and libera- 
tion of chlorine under the influence of oxygen takes place 
as represented in this equation: 

(2) 2HC1 + = H 2 + Cl 2 . 

As there is an unlimited supply of oxygen in the air, it 
would be advantageous if the decomposition of the hydro- 
chloric acid could be. effected by means of the element in 
the free state. But free oxygen alone will not accomplish 
the change. 

Deacon's Process. —A process has been invented for the 
manufacture of chlorine on the large scale which depends 
upon the decomposition of hydrochloric acid by the oxygen 
of the air. This is Deacon's process. It consists in passing 



128 COLLEGE CHEMISTRY. 

hydrochloric acid and air together through a heated tube 
containing clay balls saturated with a solution of copper 
sulphate, and then dried. If the temperature of the tube 
is not raised too high the copper sulphate remains un- 
changed. Exactly why the oxidation of the rrydrochloric 
acid takes place under these circumstances is not known, 
but it probably depends upon the formation and decom- 
position of intermediate products. Deacon's process is 
used quite extensively on the continent of Europe, while 
in England and Scotland another method, apparently more 
complicated, known as Weldon's process, is chiefly used. 

Laboratory Process. — In the laboratory chlorine is gen- 
erally made by treating hydrochloric acid with manganese 
dioxide. The reaction is represented thus: 

(1) Mn0 2 + 4HC1 = MnCl 2 + 2H 2 + Cl 2 . 

The first change should perhaps be represented thus: 

Mn0 2 + 4H01 = MnCl 4 + 2H 2 0. 

But the compound MnCl 4 gives up half its chlorine when 
heated : 

MnCl 4 = MnCl 2 + Cl 2 ; 

so that the action of hydrochloric acid on manganese 
dioxide is represented by equation (1). 

Weldon's Process. — As there is a large demand for 
chlorine, much attention has been given to the improve- 
ment of the methods for its preparation. One of the 
objections to the ordinary method is the comparatively 
high price of the mineral, manganese dioxide. As this is 
converted into the chloride, MnCl 2 , in the preparation of 
chlorine, and the chloride is of no value, the expense of 
preparation is quite high. A process has been invented 
for the regeneration of the manganous chloride, MnCl 2 , or 
for the conversion of this compound into an oxygen com- 
pound which with hydrochloric acid will give chlorine. 



PROPERTIES OF CHLORINE. 129 

This is Weldon's process. It will be taken up under the 
head of Manganese (which see). . 

Electrolytic Process. — Various processes have been 
devised for the preparation of chlorine by the action of an 
electric current on a solution of sodium chloride or of 
hydrochloric acid. Such electrolytic processes are now in 
use, and by means of them the price of chlorine has been 
much lowered. 

Properties. — Chlorine is a greenish-yellow gas. It has 
a disagreeable odor, and acts upon the membranes lining 
the throat and nose, causing irritation and inflammation, 
suggesting a " cold in the head." Inhaled in concentrated 
form it would cause death. Its specific gravity at 20° is 
2.49 (air = 1), and as compared with oxygen (16) it is 
35.45. A litre of chlorine gas, under standard conditions, 
weighs 3.22 grams. It is soluble in water and acts upon 
mercury, and therefore cannot be collected by displace- 
ment of either of these liquids. The most convenient way 
to collect it is by displacement of air. It can also be col- 
lected over warm water, in which it is less soluble than in 
cold water, or over a saturated solution of sodium chloride, 
in which it is but slightly soluble. It is easily compressed 
to a liquid and is now sold in this form in cast-iron 
cylinders. 

It is a remarkably active substance, combining with or 
acting in some way upon most other substances even at 
ordinary temperature. This activity may be illustrated 
by introducing into vessels containing chlorine a little 
finely-powdered antimony, a few pieces of thin copper-foil, 
a piece of paper with ink-marks on it, some flowers, and 
pieces of cotton-prints. The antimony will take fire and 
a white substance, will be formed. The reaction is repre- 
sented by the following equation : 

Sb + 3C1 = SbCl 3 . 

The copper-foil also takes fire, and is converted into a 
chloride as represented thus: 

Cu -f 01 = CuCl. 



130 COLLEGE CHEMISTRY. 

Many other substances unite directly with chlorine with 
evolution of heat and light, and form compounds which 
are called chlorides. This kind of action is of the same 
character as that which takes place in oxygen. There is, 
however, this difference between reactions in oxygen and 
in chlorine: the latter frequently take place at ordinary 
temperatures, whereas those in oxygen require an elevation 
of temperature to start them. In both cases the gases 
combine with the other substances directly, and disappear 
as such, and the light and heat are caused by the act of 
combination. 

Dry liquid chlorine, at its boiling temperature (— 33°. 6) 
does not act on potassium, sodium, or aluminium. 

The action of chlorine upon ink, flowers, and cotton- 
prints illustrates its power to bleach. It is important to 
notice that if the colored objects are introduced dry into 
dry chlorine the action does not take place. Moisture is 
essential to bleaching by chlorine. Chlorine acts directly 
upon some dye-stuffs, converting them into colorless sub- 
stances. In other cases it has been shown that the 
destruction of the color is due to the action of oxygen, 
which is set free from water by chlorine. In direct sun- 
light chlorine decomposes water partly according to the 
equation 

H 2 + Cl 2 = 2HCl+0; 

and partly according to the equation 

Cl 2 + H 2 = HOC1 + HC1. 

Hypochlorous 
Acid. 

This decomposition can be illustrated by filling a long tube 
with a solution of chlorine hi water, and inverting it in a 
shallow vessel containing some of the same solution. If 
this is placed in the direct sunlight, bubbles of gas will be 
seen to rise in the tube and these will collect at the top, 
while the color of the solution, which was at first greenish- 
yellow like that of chlorine, will disappear. The gas 
which collects in the upper part of the tube is oxygen. 



CHLORINE: DIFFERENT KINDS OF ACTION. 131 

The disintegrating action of chlorine upon substances 
of animal and vegetable origin may be illustrated by 
moistening a piece of filter-paper with some oil of turpen- 
tine, and introducing it into a vessel of chlorine. A flash 
of light is seen, and a dense black cloud is formed. The 
black substance is mainly carbon. Oil of turpentine is a 
compound of carbon and hydrogen. Chlorine abstracts 
the hydrogen from the carbon, leaving the latter mainly 
in the uncombined state. If the chlorine is allowed to act 
slowly upon the oil of turpentine and similar organic sub- 
stances, the chlorine is substituted atom for atom for the 
hydrogen, and a series of so-called substitution-products is 
obtained. 

Chlorine dissolves readily in water and forms a solution 
known as chlorine water. It has the odor and color of the 
gas, and it is frequently used in the laboratory instead of 
the gas. From what has been said it is evident that it 
must be kept protected from the sunlight, or decomposi- 
tion will take place, resulting in the formatiou of hydro- 
chloric acid and oxygen. 

Different Kinds of Action. — A careful study of the 
different kinds of action exhibited by chlorine shows that 
they may be classified under three heads : 

(1) First it acts by direct combination with elements as 
in the experiments with antimony and copper, and, as will 
be shown, with hydrogen and many other elements. Just 
as the compounds of oxygen with other elements are 
called oxides, so the compounds of chlorine with other 
elements are called chlorides. Thus the compound of 
antimony and chlorine, SbCl 3 , is called antimony chloride; 
that of zinc and chlorine, ZnCl 2 , is called zinc chloride; 
etc. In case an element forms more than one compound 
with chlorine, the names used to distinguish between them 
are similar to those used for oxides. Mercury forms two 
chlorides which have the composition represented by the 
formulas HgCl and HgCl 2 . The one with the smaller 
proportion of chlorine is called mercurous chloride, and the 
one with the larger proportion of chlorine is called mer- 



132 COLLEGE CHEMISTRY. 

curie chloride. So, too, there are two chlorides of iron 
which correspond to the formulas FeCl 2 and FeCl 3 . The 
former is called ferrous chloride, and the latter ferric 
chloride. 

(2) The second kind of action of chlorine is that which 
is called substitution. This was referred to in connection 
with the action of chlorine on the oil of turpentine. The 
general character of this kind of action may be illustrated 
by the following example. There is an important com- 
pound of carbon and hydrogen called benzene, which has 
the formula 6 H 6 . "When chlorine is passed through this 
compound, which is a liquid, a gas is given off which 
can easily be shown to be hydrochloric acid, HC1. This 
action continues until there is no hydrogen left in com- 
bination with the carbon, but in place of the benzene 
there is now a compound of the formula C 6 C1 6 . This has 
been shown to be the final product of a series of reactions 
represented by the following equations: 

C 6 H 6 + Cl 2 = C 6 H 5 C1 + HC1; 
C 6 H 5 C1 + 01, = C 6 H 4 C1 2 + HOI; 
C 6 H 4 C1, + Cl 2 = 6 H S 01 3 + HC1; 
C.H.C1, + 01, = C 6 H 2 C1 4 + HC1; 
C 6 H 2 C1 4 +C1 2 =C 6 HC1 5 +HC1; 
C 6 HC1 5 + Cl 2 = 6 C1 6 + HC1. 

In each stage one atom of chlorine is substituted for an 
atom of hydrogen, but the hydrogen does not escape as 
such. It combines with chlorine and passes off in the 
form of hydrochloric acid. 

(3) The third kind of action is that noticed in bleach- 
ing, which depends upon the decomposition of water and 
the escape of oxygen as already explained. This action 
does not take place in the dark, but does take place readily 
in the direct sunlight. We have, however, seen that when 
oxygen acts upon hydrochloric acid under proper condi- 
tions water is formed and chlorine set free. It appears, 



CHLORINE HYDRATE AND LIQUID CHLORINE. 133 

therefore, that, under some circumstances, this reaction 
is possible: 

2HC1 + = H 2 0-f- Cl 2 ; 

and, under other circumstances, this one : 
H 2 + Cl 2 = 2HC1 + 0. 

These facts appear to be contradictory. What part the 
sunlight plays is not known, though it is well known that 
it is capable of producing a great variety of chemical 
changes. We shall soon see that it is only necessary to 
allow it to act for an instant upon a mixture of hydrogen 
and chlorine to cause them to combine with violence. 
Then, too, the various processes known under the general 
name of photography depend upon chemical changes 
brought about by the sunlight. Leaving out of considera- 
tion this kind of action, the decomposition of hydrochloric 
acid by oxygen and that of water by chlorine can be 
explained by a consideration of the heat relations. The 
heat evolved in the formation of 1 molecule of water in the 
gaseous form is 58,069 cal., while that absorbed in the 
decomposition of 2 molecules of hydrochloric acid is 
44,002 cal. Therefore, the reaction, 

2HC1 + =- H 2 0-j-C! 2 , 

is accompanied by an evolution of heat. It is exothermic, 
and can take place without the addition of energy from 
without. If water is formed as a liquid, the heat evolved 
for 1 molecule is 68,357 cal., while that evolved by the 
formation of 2 molecules of hydrochloric acid in solution 
is 78,630 cal. Therefore, the heat evolved in the forma- 
tion of hydrochloric acid in solution is greater than that 
required to decompose water, and this reaction takes place. 
This does not, however, explain what part the sunlight 
plays in the process. 

Chlorine Hydrate and Liquid Chlorine. — When chlorine 
gas is passed into water cooled down almost to the freez- 
ing-point, crystals appear in the vessel. These consist 




134 COLLEGE CHEMISTRY. 

of chlorine and water as represented by the formnl 
CI -j- 5H 2 ; or, assuming that it is formed by the com- 
bination of the molecules of chlorine with water, the 
formula should be written Cl 2 -j- 10H 2 O. It gives off 
chlorine at the ordinary temperature and, if allowed to 
stand, undergoes complete decomposition into chlorine and 
water. If gently heated the chlorine is given off rapidly. 
This fact was taken advantage of by Faraday for the pur- 
pose of subjecting the gas to high pressure and low tern-' 
perature. For this purpose he placed some of the hydrate 

in a strong glass tube of the 
form represented in Fig. 34. 
The compound was put in the 
part of the tube marked ab, and 
the other end, c, then sealed. 
The arm ab was warmed by 
dipping it in warm water, while 
the other arm was placed in a 
freezing mixture. Under these circumstances the chlorine 
is given off from the hydrate, but being unable to escape 
from the tube the pressure is increased to such an extent 
that at the low temperature the gas assumes the liquid 
form. 

Applications of Chlorine, — Chlorine is used very exten- 
sively in the arts, particularly for the purpose of bleaching. 
It is also used in the manufacture of a large number of 
compounds which contain chlorine, the principal ones 
being bleaching powder or calcium hypochlorite, and 
potassium chlorate. If used in sufficient quantity chlorine 
is an excellent disinfectant and deodorizer. By far the 
largest quantity of the chlorine manufactured is converted 
into bleaching powder or calcium hypochlorite, as this can 
be conveniently transported, and the chlorine can be 
obtained from it when desired. It is only necessary to 
expose it to the air to effect a partial decomposition, accom- 
panied by a liberation of chlorine; and the addition of 
hydrochloric or sulphuric acid causes it to give it up com- 
pletely, as will be shown farther on. This bleaching 



HYDROCHLORIC ACID. 135 



powder is now used almost exclusively instead of cbloi 
gas for bleaching. 



me 



Hydrochloric Acid. 

Historical. — Hydrochloric acid was first prepared in 
large quantity by Glauber in the seventeenth century, and 
his description is not unlike those which one frequently 
reads nowadays- referring to some patent medicine. The 
method of preparation used by him was the same as that 
used at present, viz., the action of sulphuric acid upon 
common salt. 

Study of the Action of Hydrogen upon Chlorine. — If 
hydrogen and chlorine are brought together in the dark 
no action takes place, no matter how long they are allowed 
to stand together. If, however, the mixture is put in 
diffused sunlight, gradual combination takes place; and if 
the direct light of the sun is allowed to fall for an instant 
on the mixture, explosion occurs, and this is the sign of 
the combination of the two gases. The same sudden 
combination is effected by applying a flame or spark to the 
mixture, or by illuminating it for an instant with the 
light from burning magnesium or an electric light. On 
comparing these facts with those learned in studying the 
action of hydrogen on oxygen a marked difference is evi- 
dent. Hydrogen and oxygen do not combine either in 
the dark or the direct sunlight, but only when a spark is 
brought in contact with the mixture. 

Another way in which hydrogen can be made to com- 
bine with chlorine is by introducing a jet of burning 
hydrogen into a vessel containing chlorine. The hydrogen 
will continue to burn, but the character of the flame will 
change completely, and above the vessel white fumes will 
be observed. This burning of hydrogen in chlorine is 
entirely analogous to the burning of hydrogen in oxygen. 
It is simply an act of combination of the two gases, in each 
case accompanied by an evolution of light and heat. And 
just as oxygen can be burned in hydrogen by a proper 



136 COLLEGE CHEMISTRY. 

arrangement of apparatus, so chlorine can also be burned 
in hydrogen. 

To determine the relation between the volumes of 
hydrogen and chlorine which combine with each other and 
the volume of the product formed is more difficult than in 
the case of hydrogen and oxygen, mainly for the reason 
that chlorine acts upon mercury and is dissolved by water. ' 
It is necessary to proceed indirectly. 

Instead of causing hydrogen and chlorine to combine, 
hydrochloric acid is decomposed and the volumes of the 
hydrogen and chlorine obtained are determined. One 
method of effecting this consists in decomposing hydro- 
chloric acid by an electric current in an apparatus like 
that referred to in connection with decomposition of w T ater. 
As chlorine is, however, soluble in water, the apparatus is 
filled with a saturated solution of common salt to which a 
strong solution of hydrochloric acid in water is added. 
On passing a fairly strong current, the hydrochloric acid 
is decomposed, hydrogen being given off at one pole and 
chlorine at the other. For a given volume of hydrogen 
the same volume of chlorine is liberated, which makes it 
appear probable that hydrogen and chlorine are combined 
in hydrochloric acid in the proportion of volume to 
volume. 

For the purpose of further studying the volume rela- 
tions, the following experiment is of value: A tube is 
filled with hydrochloric acid gas. A small piece of potas- 
sium is then introduced, when decomposition takes place 
as represented in the equation 

K + HC1 --= KC1 + II. 

The gas left in the vessel is hydrogen, as can easily be 
shown. On measuring its volume it is found to be just 
half that of the hydrochloric acid gas decomposed. Taking 
this fact into consideration with the fact that whenever 
hydrochloric acid is decomposed by an electric current 
equal volumes of hydrogen and chlorine are obtained, it 
appears that in the formation of hydrochloric acid gas 



HYDROCHLORIC ACID. 137 

1 volume of hydrogen combines with 1 volume of chlorine 
to form 2 volumes of hydrochloric acid, a fact which was 
referred to in the chapter on the Atomic Theory. The 
weight of the hydrogen is found to bear to the weight of 
the hydrochloric acid the proportion 1.01:36.46. In 
other words, in 3G.4G parts by weight of hydrochloric acid 
there are 35.45 parts of chlorine and 1.01 parts of 
hydrogen. 

Preparation. — For the preparation of hydrochloric acid 
in the laboratory as well as on the large scale, common salt 
is treated with ordinary sulphuric acid. Two reactions 
may take place between these substances, depending largely 
upon the amount of sulphuric acid used. If the two sub- 
stances are brought together in the proportion of the 
weights of their molecules or their molecular weights, the 
principal reaction is the one represented in the following 
equation : 

NaCl + H 2 S0 4 = NaHS0 4 + HC1. 

Acid Sodium 
Sulphate. 

In this case the atom of sodium of the molecule of 
sodium chloride is substituted for one atom of hydrogen 
in the molecule of sulphuric acrd, while the hydrogen and 
chlorine unite to form hydrochloric acid. If, on the other 
hand, the substances are brought together in the propor- 
tion of 2 molecules of sodium chloride and 1 molecule of 
sulphuric acid the principal reaction is the following: 

2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1. 

Properties. — Hydrochloric acid is a colorless transparent 
gas, and has a sharp penetrating taste and smell. Inhaled 
it produces suffocation. It is extremely easily soluble in 
water, 1 volume of water at ordinary temperatures dissolv- 
ing -450 times its own volume of the gas, and at 0°, 500 
times. The solution of the gas in water is what is gen- 
erally called hydrochloric acid. So great is the attraction 
of the gas for water that it condenses moisture from the 
air; hence, although the gas itself is quite colorless and 



13^ COLLEGE CHEMISTRY. 

transparent, when it comes in contact with the air dense 
white clouds are formed, which are not formed if it is kept 
from contact with the air, as can easily be shown by filling 
glass vessels with the gas. Hydrochloric acid does not 
burn and does not support combustion. This is equivalent 
to saying that it does not combine with oxygen under 
ordinary circumstances, and that substances which com- 
bine with the oxygen of the air do not combine with 
hydrochloric acid. On the other hand, we have seen that 
under some circumstances oxygen does act upon hydro- 
chloric acid and cause an evolution of chlorine. The gas 
is comparatively easily condensed to the liquid form. 

When a concentrated solution of hydrochloric acid in 
water is heated, gas is given off, but if a dilute solution is 
heated, water is given off. In either case, when the com- 
position of the liquid is that represented by the formula 
HC1 -|- 8H 2 0, it boils under the ordinary pressure of the 
atmosphere unchanged. If the pressure is lowered the 
composition of the liquid which passes over in the process 
of distillation changes, so that it contains a larger per- 
centage of hydrochloric acid the lower the pressure 
becomes. This fact seems to show that the liquid of the 
composition HC1 -f- 8H 2 0, which boils unchanged at the 
temperature 110° under the ordinary pressure of the 
atmosphere, is not a chemical compound. On the other 
hand, it does not conduct itself like most ordinary 
solutions of gases. 

There is a definite compound of hydrochloric acid with 
water called hydrochloric acid hydrate, which has the com- 
position HC1 -f- 2H 2 0. This is formed by passing hydro- 
chloric acid gas into the concentrated aqueous solution 
cooled down to — 22°. Under these circumstances the 
hydrate separates in the form of crystals. 

Commercial hydrochloric acid is a yellowish liquid, the 
color being due to the presence of impurities, such as iron 
and organic substances. The solution is obtained in the 
factories in which "soda" or sodium carbonate is made. 
This is an extremely important substance in the arts. It 



HYDROCHLORIC ACID. 139 

does not occur in nature, but is manufactured from com- 
mon salt. In the process most commonly used salt is first 
converted into sodium sulphate, Na 2 S0 4 , by treating it 
with sulphuric acid. Hydrochloric acid is necessarily 
given off. When the factories were first established in 
England, the gas was allowed to escape as a waste product, 
but the effects produced by it upon the vegetation of the 
surrounding country were so injurious that a law was 
passed prohibiting the manufacturers from allowing the 
gas to escape. It is now collected by causing it to pass 
through towers so constructed as to expose a large surface 
of bricks or sandstone plates over which a current of cold 
water is constantly kept flowing. This water dissolves the 
hydrochloric acid, and the solution collected below is com- 
mercial hydrochloric acid. In this way enormous quanti- 
ties of the acid are produced, but its uses are numerous 
and it always commands a price. 

Pure hydrochloric acid is a solution of the pure gas in 
pure water. It is colorless, and when concentrated it gives 
off fumes when exposed to the air. The solution when 
heated gives off a large part of the gas contained in it, and 
by boiling it can all be evaporated. 

Chemical Action of Hydrochloric Acid. — If the action 
of hydrochloric acid towards the elements should be studied 
systematically it would be found that many of them act 
by simply taking the place of the hydrogen, as has already 
been illustrated in the preparation of hydrogen by treating 
hydrochloric acid with zinc, when this reaction takes place : 

Zn + 2HC1 = ZnCl 2 + H 2 . 
With iron the reaction is : 

Fe + 2HC1 = FeCl 2 + H 2 ; 
with tin : 

Sn + 2HC1 = SnCl 2 + H t ; 

with potassium : 

K + HC1 = KC1 -f H, or 
K 2 + 2HC1 = 2KC1 + H 2 ; 



14° COLLEGE CHEMISTRY. 

with sodium: 

Na 2 -f-2HCl = 2NaCl + H 2 ; 
with calcium : 

Ca -f 2HC1 = CaCl 2 + H 2 ; etc., etc. 

On the other hand, there are many elements that do not 
act in this way towards hydrochloric acid. Sulphur, 
nitrogen, phosphorus, carbon, and boron may be taken as 
examples. These elements do not act upon hydrochloric 
acid at all. We might, therefore, divide the elements into 
two classes: (1) those which act upon hydrochloric acid 
setting hydrogen free and forming chlorides; and (2) those 
which do not act upon hydrochloric acid. 

Again, when hydrochloric acid acts upon the oxides of 
those elements which have the power to liberate hydrogen 
from it, it forms the same chlorides as are formed when 
the element alone acts, but instead of hydrogen being 
liberated water is formed. Thus when the acid acts upon 
zinc oxide, ZnO, the reaction takes place thus: 

ZnO -f 2HC1 = Zn01 2 + H 2 0; 

with lime or calcium oxide, CaO, it is : 

CaO-f 2HCl = CaCl 8 + H 2 0; 

with potassium oxide: 

K 2 + 2HC1 = 2KC1 -f H 2 0: etc., etc. 

When the so-called hydroxides of those elements which 
act upon hydrochloric acid are brought in contact with the 
acid, action takes place even more readily than with the 
oxides, and the products are the same. Thus potassium 
hydroxide and hydrochloric acid give potassium chloride 
and water: 

KOH + HC1 = KC1 + H 2 0; 

calcium hydroxide and hydrochloric acid give calcium 
chloride and water, thus: 

Ca(OH) 2 + 2HC1 = CaCl 2 + 2H 2 0; 



HYDROCHLORIC ACID. 141 

aluminium hydroxide and hydrochloric acid give alumin- 
ium chloride and water, thus: 

A1(0H) 8 + 3H01 = A101, + 3H 2 0; etc., etc. 

There are then elements which act upon hydrochloric 
acid liberating hydrogen and forming chlorides; and the 
oxides and hydroxides of these elements act upon hydro- 
chloric acid forming chlorides and water. The elements 
which act in this way are commonly called metals or, for 
reasons which will be discussed farther on, base-forming 
elements. 

If those elements which do not set hydrogen free from 
hydrochloric acid are treated directly with chlorine, they 
generally combine with it to form chlorides. But these 
chlorides differ markedly from the chlorides of the metals, 
especially in their conduct towards water. Two examples 
will suffice for the present. Phosphorus forms a chloride 
of the formula PC1 3 , known as phosphorus trichloride. 
In contact with water it undergoes decomposition accord- 
ing to this equation : 

PCI3 + 3H 2 = P0 3 H 3 + 3HC1; 

so, too, the chloride of boron, BC1 3 , undergoes the same 
kind of change: 

BC1 3 + 3H 2 = B0 3 H 3 + 3HC1. 

Similarly, the other chlorides of the elements of this 
class tend to pass into the oxides or hydroxides when 
brought in contact with water. Those elements which do 
not act upon hydrochloric acid setting hydrogen free and 
forming chlorides are generally called non-metals or, for 
reasons which will appear later, acid-forming elements. 
The chlorides of the acid- forming elements are generally 
decomposed by water and the corresponding oxides or 
acids are formed. In general terms, the oxide of a base- 
forming element or of a metal is acted upon by hydro- 
chloric acid, a chloride and water being formed; and the 



142 



COLLEGE CHEMISTRY. 



chloride of an acid-forming element or of a non-metal is 
acted upon by water, an acid, or oxide, and hydrochloric 
acid being formed. We shall have many illustrations of 
the opposite chemical character of these two classes of 
elements, and we shall see that many of the most important 
and characteristic chemical reactions are associated with 
these differences. 



EXPERIMENTS. 

Experiment 65. — Pour 2 or 3 cc, concentrated sulphuric acid 
on a gram or two of common salt in a test-tube. A gas will be 
given off which forms dense white fumes in the air and has a 
sharp, penetrating taste and smell. This is hydrochloric acid. 

Experiment 66. — The best way to make chlorine is the follow- 
ing : Mix 5 parts (say 50 grams) coarsely-granulated manganese 
dioxide and 5 parts coarsely-granulated common salt. Make a 

mixture of 12 parts concen- 
trated sulphuric acid and 6 
parts water. Let this mixture 
cool down to the temperature 
of the room, and then pour it 
upon the mixture of salt and 
manganese dioxide. Gently 
heat on a sand-bath, and a 
regular current of chlorine 
will be given off. The gas is 
collected by displacement of 
air in a dry glass vessel. The 
apparatus for the purpose is 
arranged as shown in Fig. 35. 
The delivery-tube should reach 
to the bottom of the collecting 
vessel, and the mouth of the 
vessel should be covered with 
a piece of paper to prevent currents of air from carrying away 
the chlorine. As the gas collects in the vessel the experimenter 
can judge of the quantity present by means of the color. 

Experiment 67. — Collect six or eight dry cylinders or bottles 
full of chlorine. Make the gas from about 30 grams of manga- 
nese dioxide, using the other substances in the proportion al- 
ready stated. 




Fig. 35. 



- 



CHLORINE: EXPERIMENTS. 



*43 



(1) Introduce into one of the vessels containing chlorine a 
little finely-powdered antimony. 

(2) Into a second vessel put a few pieces of heated thin copper- 
foil. 

(3) Into a third vessel put a piece of paper with some writing 
on it, some flowers, and pieces of cotton print. The substances 
used must be moist. 

(4) Into a fourth vessel put a dry piece of the same cotton 
print as that used in the previous experiment. 

What conclusions do the results of the above experiments 
justify as to the conduct of chlorine ? 

Experiment 68.— Cut a piece of filter-paper about an inch wide 
and six to eight inches long. Pour on this some ordinary oil of 
turpentine previously warmed slightly. Introduce this into one 
of the vessels of chlorine. A flash of flame is noticed, and a 
dense black cloud is formed. The action in this case is due to 
the great affinity of chlorine for hydrogen. Oil of turpentine 
cousists of carbon and hydrogen. The main action of the chlo- 
rine consists in extracting the hydrogen and leaving the carbon. 
The experiment is interesting chiefly in so far as it illustrates the 
geueral tendency of chlorine to act upou vegetable substances. 

Chlorine decomposes Water in the Sunlight. 



Experiment 69.— Seal the end of a glass tube about a metre 
(or ;ibout a yard) long and about 12 mm. (i inch) 
internal diameter. Fill this with a strong solu- 
tion of chlorine in water. Invert it as shown in 
Fig. 36, in a shallow vessel containing some of 
the same solution of chlorine in water. Place 
the tube in direct sunlight. Gradually bubbles 
of gas will be seen to rise and collect in the upper 
end, and the color of the solution, which is at 
first greenish yellow, like that of chlorine, dis- 
appears. The gas can be shown to be oxygen. 

Chlorine Hydrate. 

Experiment 70. — Conduct chlorine into a flask 
containing water cooled down to about 2° or 3° 
C. If crystals are formed remove some by filter- 
ing out-of-doors if the weather is cold. Expose 
some of the crystals on filter-paper under a hood 
in the laboratory. What changes have taken place ? 



Fig. 36. 



44 



COLLEGE CHEMISTRY. 



Formation of Hydrochloric Acid. 

Experiment 71. — Light a jet of hydrogen in the air and care- 
fully introduce it into a vessel containing chlorine. It will con- 
tinue to burn, but the flame will not appear the same. A gas 
will be given off which forms clouds in the air. This gas lias a 
sharp, penetrating taste and smell. 

Experiment 72. — Half fill a small, wide-mouthed cylinder over 
salt water with chlorine gas. Then fill it with hydrogen. The 
direct sunlight must not shine upon the cylinder while it contains 
the mixture. Turn it mouth upward and apply a flame. 

Preparation of Hydrochloric Acid. 
Experiment 73. — Arrange an apparatus as shown in Fig. 37. 




Fig. 37. 

Weigh out 5 parts common salt, 5 parts concentrated sul- 
phuric acid, and 1 part water. Mix the acid and water, taking 
the usual precautions ; let the mixture cool down to the ordinary 
temperature, and then pour it on the salt in the flask. For the 
purposes of the experiment take about 20 grams of salt. Now 
heat the flask gently, and the gas will be regularly evolved. Con- 
duct it at first over water contained in the two Woulff 's bottles. 
When the air has been displaced, the gas is all absorbed as soon as 
it comes in contact with the water. — After the gas has passed for 
ten to fifteen minutes, disconnect at A. Notice the fumes. These 
become denser by blowing the breath on them. Why ?— Apply a 
lighted match to the end of the tube. Does the gas burn?— 



HYDROCHLORIC ACID: EXPERIMENTS. 1 45 

Collect some cf the gas in a dry cylinder by displacement of air, 
;>s in the case of chlorine. The specific gravity of the gas being 
1.26, the vessel must of course be placed with the month upward. 
That the gas is colorless and transparent is shown by the appear- 
ance of the generating- flask, which is filled with the gas. Insert a 
burning stick or candle in the cylinder filled with the gas.— Re- 
connect the generating flask with the series of bottles containing 
water, and let the process continue until no more gas comes over. 
The reaction represented in the equation 

2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1 

is now complete. Disconnect the flask, and after it has cooled 
down pour water on the contents ; when the substance is dissolved 
filter it and evaporate to such a concentration that, on cooling, 
the sodium sulphate is deposited. Pour off the liquid and dry 
the solid substance by means of filter-paper. Compare the sub- 
stance with the common salt which you put in the flask before the 
experiment. What proofs have you that the two substances are 
not the same ?— Heat a small piece of each in a dry tube closed at 
one end. What differences do you notice ? — Treat a small piece 
of each in a test-tube with sulphuric acid. What difference do 
you notice ? — If in the experiment w r e should recover all the sodium 
sulphate formed, how much should we have ?— Put about 50 cc. 
of the liquid from the first Woulff's bottle in a porcelain evapora- 
ting-dish. Heat over a small flame just to boiling. Is hydro- 
chloric acid given off ? Can all the liquid be driven off by boil- 
ing ? — Try the action of the solution on some iron filings. What 
is given off ?— Add some to a little granulated zinc in a test-tube. 
What is given off ? — Add a little to some manganese dioxide in a 
test-tube. What is given off ? — Add ten or twelve drops of the 
acid to 2 to 3 cc. water in a test-tube. Taste the dilute solution. It 
has what is called a sour or acid taste, the two terms being practi- 
cally synonymous. — Add a drop or two of a solution of blue litmus, 
or put into it a piece of paper colored blue with litmus. What 
change takes place ? Litmus is a vegetable color prepared for use as 
a dye. Other vegetable colors are changed by hydrochloric acid. — 
The color will be restored by adding a few 7 drops of a solution of 
caustic soda.— In what experiment has caustic soda been obtained ? 
What relation does it bear to water? — To the dilute solution of 
hydrochloric acid add drop by drop a dilute solution of caustic 
soda. Is the acid taste destroyed ? 



CHAPTER IX. 

COMPOUNDS OF CHLORINE WITH OXYGEN AND 
WITH HYDROGEN AND OXYGEN. 

General. — As has been seen, chlorine combines very 
readily with hydrogen, and hydrogen with oxygen, and the 
products are stable compounds. On the other hand, 
chlorine cannot be made to combine directly with oxygen. 
By indirect processes they can be combined, but the com- 
pounds undergo decomposition easily, yielding back the 
chlorine and oxygen contained in them. Before taking 
up the compounds of chlorine and oxygen, however, it will 
be best to discuss, as far as may be necessary, the com- 
pounds of chlorine, hydrogen, and oxygen which are more 
easily made, and from which the oxides are made. 

Principal Reactions for making Compounds of Chlorine 
with Hydrogen and Oxygen. — When chlorine acts upon 
water in the direct sunlight one of the reactions that takes 
place is represented by the equation 

Cl 2 + H 2 = HOC1 -f- HCL 

The compound represented by the formula HOC1 is called 
hvpochlorous acid. So, also, when chlorine acts upon a 
dilute solution of potassium hydroxide, KOH, a similar 
reaction takes place. It is represented thus : 

Cl 2 + KOH = KOC1 +HC1. 

The hydrochloric acid formed acts upon some of the 
potassium hydroxide thus : 

KOH + HC1 = KC1 + H 2 0. 

146 



COMPOUNDS OF CHLORINE. 147 

Or, combining the two equations we have: 

Cl 2 + 2KOH = KOC1 + KC1 + II 2 0. 

The product represented by the formula KOC1 is known 
as potassium hypochlorite, and that represented by the 
formula KC1 is known as potassium chloride. 

When the temperature of a solution of potassium hypo- 
chlorite is heated, the substance undergoes a change as 
represented thus : 

3KOC1 = KCIO, + 2KC1. 

The compound represented by the formula KCIO, is 
known as potassium chlorate. This has already been em- 
ployed in the preparation of oxygen. 

When, further, chlorine is passed into a warm concen- 
trated solution of potassium hydroxide all the above reac- 
tions take place, and they may be represented in one 
equation : 

6KOH + 3C1 2 = 5KC1 + KC10 3 + 3H 2 0. 

Potassium chlorate. KC10 3 , and potassium hypochlorite, 
KCIO, bear the same relation to two compounds, HC10 3 
and HCIO, that potassium chloride, KC1, and sodium 
chloride, NaCl, bear to hydrochloric acid. But we have 
seen that hydrochloric acid can easily be obtained from 
sodium chloride by treating it with sulphuric acid. Potas- 
sium chloride undergoes the same change when treated 
with sulphuric acid. Indeed, we shall see that nearly all 
compounds containing sodium or potassium give up these 
metals when treated with sulphuric acid, and take up 
hydrogen in the place of them. 

Treating potassium chloride with sulphuric acid this 
reaction takes place : 

2KC1 -f H 2 S0 4 = K 2 S0 4 + 2HC1. 

Similarly, treating potassium chlorate with sulphuric acid, 
this reaction takes place : 

2KC10 3 + H 2 S0 4 = K 2 S0 4 + 2HC10 3 . 



148 COLLEGE CHEMISTRY. 

The compound HC10 3 is called cldoric acid. Further, 
when potassium hypochlorite is decomposed by sulphuric 
acid under proper circumstances, it undergoes the same 
kind of decomposition : 

2KC10 + H 2 SO, - K 2 S0 4 + 2HC10. 

The compound HCIO is called hypochlorous acid. 

Hypochlorous Acid, HCIO. — The formation of this com- 
pound by the action of chlorine on water and that of 
potassium hypochlorite, KCIO, by the action of chlorine 
oh caustic potash have been mentioned. A similar reaction 
is employed on the large scale in the manufacture of 
bleaching powder or " chloride of lime." This consists in 
treating slaked lime or calcium hydroxide with chlorine. 
The action is represented thus : 

2Ca(OH) 2 + 201 2 = Ca(C10) 2 + Ca01 2 + 2H 2 0. 

The compound Ca(010) 2 , known as calcium hypochlorite, 
is derived from hypochlorous acid by substituting one 
atom of the bivalent metal calcium for two atoms of 
hydrogen in two molecules of the acid : 

HCIO . n /C10\ n /rnm 
HCIO glves 0a vcioj or Ca ( C10 )2- 

The concentrated solution of hypochlorous acid has a 
peculiar odor suggesting that of chlorine. It is the odor 
which is familiar as that of bleaching powder or chloride 
of lime. The acid undergoes decomposition very readily, 
forming chlorine and a compound of chlorine and oxygen. 
A solution of the acid bleaches about as well as chlorine, 
and when bleaching powder is used for bleaching it is 
largely the hypochlorous acid set free from the hypochlorite 
which effects the desired changes. 

Like chloric acid, hypochlorous acid is an excellent 
oxidizing agent, and is used in the laboratory in this 
capacity. 



COMPOUNDS OF CHLORINE. H9 

Perchloric Acid, HC10 4 . — When the preparation of 

oxygen by heating potassium chlorate was considered, it 
was pointed out that in the first stage of the decomposition 
a reaction of this kind t&kes place: 

2KC10 3 = KC1 + KC10 4 + 2 . 

The compound KC10 4 , or potassium perchlorate, can be 
separated from the chloride by treating the mixture with 
cold water in which the chloride is easily soluble, while 
the perchlorate is practically insoluble. When the per- 
chlorate is treated with sulphuric acid, perchloric acid can 
be obtained from the mixture by distillation. 

Pure perchloric acid, HC10 4 , can be obtained in the 
form of a colorless fuming liquid. It is a dangerous sub- 
stance to deal with, as it produces bad wounds when 
brought in contact with the flesh, and is quite unstable. 
In contact with combustible substances in general it causes 
explosion in consequence of the ease with which it gives 
up oxygen and converts the combustible substances into 
gaseous products. 

General. — From the above it will be seen that the com- 
pounds of chlorine with hydrogen and oxygen form a 
series, the members of which bear a simple relation to one 
another. Beginning with hydrochloric acid the series is 
as follows : 

Hydrochloric acid HC1 

Hypochlorous acid HCIO 

Chlorous acid HC10 2 

Chloric acid HC10 3 

Perchloric acid HCIO, 

The successive members differ from each other by one atom 
of oxygen, the ratio between the hydrogen and the chlorine 
remaining the same throughout the series. While the 
compounds differ markedly from one another in many 
ways, they have some common features. Upon metals, 
and their oxides and hydroxides, all the members of the 



150 COLLEGE CHEMISTRY. 

series act in general in the same way that hydrochloric 
acid does, the result being the formation of products 
which do not contain hydrogen, but do contain a metal in 
the place of the hydrogen. We have examples of these 
compounds in potassium chlorate, KC10 3 , calcium hypo- 
chlorite, Oa(C10) 2 , potassium chlorite, KC10, , and potas- 
sium perchlorate, KCIO^. All these compounds belong 
to the class called salts, which will presently be taken up. 
On the other hand, while there is a class of elements upon 
which hydrochloric acid does not act, the oxygen com- 
pounds of the above series w r ill in many cases act upon 
these elements and convert them into oxides. Thus sul- 
phur and phosphorus, which are not acted on by hydro- 
chloric acid, are converted into oxides by the oxygen 
compounds of chlorine. 

Finally, the addition of oxygen to hydrogen and chlorine 
decreases the stability of the compound. Hydrochloric 
acid, for example, is characterized by great stability, while 
hypochlorous acid, II CIO, as well as all the other members 
of the series, is characterized by instability. The larger 
the proportion of oxygen, however, the greater the stability 
of the compound. The most stable member of the series 
of oxygen compounds is perchloric acid. Another fact 
that is worthy of special notice is that the metal derivatives 
or salts of these acids are more stable than the acids them- 
selves. Many of them can be heated to a comparatively 
high temperature without undergoing decomposition. 
This is most marked in the case of the perchl orates. It 
will be remembered that in decomposing potassium chlorate 
for the purpose of making oxygen the change takes place 
in two stages. In the first, potassium perchlorate is 
formed. In order to decompose this, however, the tem- 
perature must be raised considerably higher than that 
which was required to effect the breaking down of the 
chlorate. 

Compounds of Chlorine with Oxygen. — The compounds 
of chlorine with oxygen are : 

Chlorine monoxide, C1 2 0, chlorine dioxide, CIO,, and 



EXPERIMENTS WITH COMPOUNDS OF CHLORINE. 15 1 

chlorine hept oxide, C1 2 7 . These ;irc unstable substances 
that easily break down into chlorine and oxygen. 



EXPERIMENTS. 

Chloric Acid and Potassium Chlorate. 

Experiment 74. — Dissolve 40 grams (about 1\ ounces) potas- 
sium hydroxide in 100 ce. water in a beaker-glass, and pass into 
it chlorine, generated as in Experiment 67, using 75 grams salt 
and the other reagents in proportion. Arrange an inverted fun- 
nel on the end of the delivery-tube, so that the edge of the funnel 
dips just below the surface of the solution in the beaker-glass, to 
prevent the choking of the delivery-tube. When the solution no 
longer shows an alkaline reaction stop the current of chlorine, and 
boil the solution for a few minutes. When the solution cools, 
crystals of potassium chlorate mixed with a little potassium chlor- 
ide will be deposited. Filter these off, dissolve in a small volume 
of water, filter, and allow to cool. If the solution is sufficiently 
concentrated crystals will be deposited. Filter off the crystals 
and dry them. What evidence have you that the substance is 
potassium chlorate ? Does it give off oxygen when heated ? In 
a dry test-tube pour two or three drops of concentrated sulphuric 

acid on a small crystal of 
the substance. Do the same 
with a piece of potassium 
chlorate from the labora- 
tory bottle. Hold the 
mouth of the test-tube 
away from the face. What 
is noticed in each case ? — 
Evaporate the solution 
from which the crystals of 
potassium chlorate have 
been removed. On allow- 
ing it to cool crystals will 
again be deposited. Take 
them out and recrystallize 
them. Does this sub- 
stance give off oxygen 
when heated ? Does it give 
off a gas when treated with sulphuric acid ? Is this gas colored \ 
Is it hydrochloric acid? How do you know that it is? If the 




Fig 



15^ 



COLLEGE CHEMISTRY. 



gas is hydrochloric acid, what is the solid substance from which 
it is formed ? And what is left in the test-tube ? 

Experiment 75.— Place the slaked lime, made by slaking 20 to 
30 grams of quicklime with enough hot water to make a dry 
powder, in a 250-cc. flask with a two-hole stopper. Introduce a 
glass tube through one of the holes and pass chlorine in slowly, 
while constantly shaking the flask, for about ten minutes. Intro- 
duce a funnel and delivery-tube as showm in Fig. 38. Pour a 
mixture of equal parts of sulphuric acid and water slow r ly through 
the funnel-tube. Collect by displacement of air the gas given 
off. What evidence have you that the gas is chlorine ? 

Perchloric Acid. 



Experiment 76. — Make potassium perchlorate as follows: 
Gently heat 50 to 100 grams potassium chlorate until after having 
been liquid it becomes thick and pasty, and gas is not given off 
without raising the temperature. After cooling, break up the 
mass and treat it with cold water. This dissolves out the potas- 
sium chloride and leaves the perchlorate, which can then be crys^ 
tallized from hot w T ater. After the crystallized salt is dried it is 
decomposed by sulphuric acid. To effect this decomposition, the 
finely-powdered salt (10 parts) is treated in a retort with 20 parrs 
of pure sulphuric acid which is free from nitric acid and diluted 
with -jV its volume of water. The retort is connected with a 
receiver which can be well cooled. The mixture is heated, and 
when the perchloric acid begins to come over, the heat is so regu- 
lated that the temperature does not rise above 140°. When the 
mixture has become colorless the operation is ended. 



CHAPTER X. 
ACIDS.— BASES.— NEUTRALIZATION.— SALTS. 

General. — One cannot deal with chemical phenomena 
without constant reference to acids, and in the course of 
our study thus far a number of substances belonging to 
this class have been met with. It is now time to inquire 
what features these substances have in common which lead 
chemists to call them all acids. What is there in common 
between the heavy, oily liquid, sulphuric acid, the colorless 
gas, hydrochloric acid, and the unstable compounds chloric 
and hydrochlorous acids ? To understand the common 
features requires some knowledge of a class of substances 
to which attention has already been given. These are 
substances like caustic potash and caustic soda, or potas- 
sium and sodium hydroxides, called alkalies, which are the 
most marked representatives of the class of substances 
known as oases. These two classes, acids and bases, have 
the power to destroy the characteristic properties of each 
other. When an acid is brought in contact with a base in 
proper proportions, the characteristic properties of both 
the acid and the base are destroyed. They are said to 
neutralize each other. They form new products which 
are said to be neutral, which means that they have not the 
properties of an acid nor those of a base. This act of 
neutralization is an extremely important one, with which 
we have constantly to deal in chemical operations. 

A Study of the Act of Neutralization.— The fact having 
been learned that acids and bases neutralize one another, 
the next thing to do is to study the act of neutralization 

153 



154 COLLEGE CHEMISTRY. 

as carefully as, possible, and learn what chemical changes 
are involved in it. For this purpose we should select a 
number of acids and a number of bases and study their 
action upon one another. We may take sulphuric, hydro- 
chloric, and nitric acids; and potassium, sodium, and 
calcium hydroxides. We know from many analyses that 
have been made that the composition of these substances 
is as follows : 

Hydrochloric acid HC1 

Nitric acid HN0 3 

Sulphuric acid H 2 S0 4 



Potassium hydroxide KOH f - 

Sodium hydroxide NaOH 

Calcium hydroxide Ca(OH) 2 ^ 

The first question to be answered is whether, in order 
to effect neutralization, definite quantities of the substances 
are necessary. To decide this, solutions of the acids and 
of the bases should be prepared and allowed to act upon 
one another in different proportions. But how shall we 
determine whether the solutions we are working with are 
acid, basic, or neutral ? It has been found that all acids 
have the power to change the color of certain substances. 
For example, the dye litmus is blue. If a solution which 
is colored blue with litmus is treated with a drop or two 
of an acid, the color is changed to red. If now the red 
solution is treated with a few drops of a solution of a 
strong base, the blue color is restored. There are many 
other substances which change markedly in color by the 
addition of acids or bases. These facts furnish a means 
of recognizing whether a solution is acid or basic. Now, 
suppose that to a carefully-measured quantity of one of 
the acid solutions a few drops of blue litmus is added. It 
will at once turn red. On adding slowly a solution of one 
of the bases the color will remain red as long as the solu- 
tion is acid, but the instant it is basic it will turn blue. 
By noticing when the change in color takes place, it is 



NEUTRALIZATION. 155 

possible to determine exactly how much of a certain basic 
solution is required to neutralize the quantity of the acid 
solution taken. If it is found in the case studied that to 
neutralize 20 cc. of the acid solution 30 cc. of the basic 
solution are required, then, using the same solutions, it 
will be found in every experiment that the same quantities 
are required to effect neutralization, or that the change of 
color takes place whenever these proportions are reached. 
And no matter how the quantity of one of the liquids is 
varied, tbe quantity of the other required for neutraliza- 
tion varies in the same proportion. A great many experi- 
ments of this kind have been performed with many different 
acids, and what is true in one case has been found true in 
all. It appears, therefore, that the act of neutralization is 
a definite one, which takes place between definite quantities 
of acid and base; that for a certain quantity of base a cer- 
tain quantity of acid is required to effect neutralization, 
and vice versa. 

The next question to be answered is, What is formed 
when the acid and base are neutralized ? To determine 
this, larger quantities of acids should be neutralized with 
bases, and the substance or substances formed should then 
be studied. If hydrochloric acid is neutralized with 
sodium hydroxide a solid product, sodium chloride, is 
formed. The action takes place according to the following 
equation : , 

HC1 + NaOH = NaOl + H*0. 

Hydrochloric acid and calcium hydroxide act thus : 
2HC1 + Ca(OH) 2 = CaCl 2 + 2H 2 0. 

Nitric acid acts upon the three bases mentioned above as 
represented in these equations : 

HN0 3 +KOH =KN0 3 +H 2 0; 
HNO, +NaOH = NaNO, + H 2 0; 
2HNO3 -f Ca(OH) 2 = Ca(NO s ), -f H 2 0. 



*56 COLLEGE CHEMISTRY. 

Sulphuric acid acts upon these same bases thus: 

H 2 S0 4 + 2KOH = K 2 S0 4 + 2H 2 0; 
H 2 S0 4 + 2NaOH = Na 2 S0 4 + 2H 2 0; 
H 2 S0 4 + Ca(OH) 2 = CaS0 4 + 2H 2 0. 

The reactions which take place in these cases are typical 
of all reactions between acids and bases. One of the 
products formed is always water, the other is a compound 
without acid and basic properties. It is neutral and differs 
from the acid in that it contains some other element in 
place of the hydrogen. This other element is the one 
which in the base is in combination with hydrogen and 
oxygen as a hydroxide. 

General Statements. — Considering the facts treated of 
in the last paragraph, it appears : 

(1) That an acid contains hydrogen; 

(2) That a base contains a metal; 

(3) That when an acid acts upon a base the hydrogen 
and metal exchange places; 

(4) That the substance formed by substituting hydrogen 
for the metal of the base is water; 

(5) That the substance obtained from the acid by sub- 
stituting a metal for the hydrogen is neither an acid nor a 
base, but is generally neutral. 

The last statement is subject to some modification, for 
reasons which in some cases are clear but in others are not 
apparent. It is true that in some cases after substituting 
a metal for the hydrogen the substance has an alkaline 
reaction, and in other cases an acid reaction. 

Definitions. — We have already seen that hydrochloric 
acid and sulphuric acid act upon certain metals, as iron 
and zinc, and that the action consists in giving up 
hydrogen and taking up metal in its place. The products 
of this action are the same in character as those formed 
by the action of acids on bases. 

An acid is a substance containing hydrogen, which 
it easily exchanges for a metal, when treated with a 



DISTINCTION BET W E EN ACIDS AND BASE'S. 157 

metal itself, or with a compound of a metal, called a 
base. 

A base is a substance containing a metal combined with 
hydrogen and oxygen (hydroxyl). It easily exchanges its 
metal for hydrogen when treated with an acid. 

The products of the action of an acid on a base are, 
first, water, and, second, a neutral substance called a salt. 

In the examples above cited the products KN0 3 , potas- 
sium nitrate; NaN0 3 , sodium nitrate; Ca(NO s ) 2 , calcium 
nitrate; K„S0 4 , potassium sulphate; Na 2 S0 4 , sodium sul- 
phate; CaS0 4 , calcium sulphate, are salts. The relations 
between them and the acids from which they are derived 
will be easily recognized on comparing their formulas with 
those of the acids. 

Distinction between Acids and Bases. — Although there 
is no difficulty in distinguishing between most acids and 
most bases, there are some compounds which act some- 
times in one way and sometimes in the other. Sulphuric 
acid, nitric acid, and hydrochloric acid always act as acids, 
and sodium and potassium hydroxides always act as bases, 
but some substances which are generally basic will under 
some circumstances act as acids, and some which act as 
acids will occasionally act as bases. What is the standard ? 
How shall we tell whether a substance is an acid or a 
base ? We may take a pronounced acid, such as hydro- 
chloric acid, and say that any hydroxide which has the 
power to neutralize this acid and form with it a salt shall 
be called a base; and in the same way we may take a 
pronounced base, like potassium hydroxide, and say that 
any hydroxide which has the power to neutralize this shall 
be called an acid. Having made the division in this way, 
it would be found that a few substances would be included 
-in both lists, or, in other words, some substances which 
are basic toward hydrochloric acid are acid toward potas- 
sium hydroxide. As an example, we may take aluminium 
hydroxide, Al(OH) 3 . This neutralizes hydrochloric acid 
and forms aluminium chloride according to the equation 

Al(OH) 3 -f- 3HC1 = AICI3 + 3H 2 0. 



X5 8 COLLEGE CHEMISTRY. 

But it also neutralizes potassium hydroxide according to 
the equation 

Al(OH), + 3KOH = Al(OK) 3 + 3H 2 0. 

It may be said in regard to this case, as in regard to most 
other cases of the kind, that the hydroxide in question is 
basic toward nearly all substances toward which potassium 
hydroxide is basic; whereas it is acid toward only three or 
four of the most energetic bases. Bearing in mind, then, 
the fact that there are some exceptional cases, it may be 
said that the distinction between acids and bases is easily 
recognized. 

Metals or Base-forming Elements. — The question, What 
is a metal ? may fairly be asked. But unfortunately it is 
by no means an easy matter to give a satisfactory answer 
to the question. We can give examples of metals, such 
as iron, zinc, silver, calcium, magnesium, etc. ; but when 
we attempt to find the distinguishing features of these 
substances we are somewhat at a loss to state them. In 
general, it may be said that to the chemist any element is 
a metal which with hydrogen and oxygen forms a base or 
a product which has the power to neutralize acids. In 
general, any element which has the power to enter into an 
acid in the place of the hydrogen is called a metal, or 
is said to have metallic properties. This is the sense in 
which the word metal is used in this book. A better, 
though a longer, name for the metals is base-forming 
elements. 

Constitution of Acids and Bases. — As has been pointed 
out, the bases are hydroxides, and these hydroxides are 
regarded as derived from water by the substitution of 
metals for the hydrogen. Examples of the hydroxides of 
univalent, bivalent, and trivalent metals were given in 
a previous chapter (see p. 111). Similarly, the acids 
which contain oxygen are regarded as hydroxides, or as 
derived from water. This view is illustrated by the fol- 
lowing formulas of some of the more common acids: 



CONSTITUTION OF ACIDS AND BASES. 159 

Nitric acid (HO)NO, 

Sulphuric acid (1I0) 2 S0 2 

Phosphoric acid (HO) 3 PO 

■Carbonic acid (HO) 2 CO 

Metaphosphoric acid. (HO)P0 2 

Nitrous acid (HO)NO 

Arsenious acid (HO) 3 As 

Hypochlorous acid (HO)Cl 

Perchloric acid (HO)C10 3 

There are three classes of acids represented in this list : 
(1) those with one, (2) those with two, and (3) those with 
three atoms of hydrogen in the molecule. Or, considering 
the compounds as hydroxides, these classes are: (1) those 
derived from one molecule, (2) those derived from two 
molecules, and (3) those derived from three molecules of 
water by substitution of something else for half the 
hydrogen. It is interesting to observe, also, that this 
something which is substituted for the hydrogen is in 
most cases an element in combination with oxygen or, if 
it is not in combination with oxygen, it has the power to 
take up more oxygen. Thus hypochlorous and arsenious 
acids are regarded as derived from water by the substitu- 
tion of chlorine and arsenic for hydrogen in water as shown 

H-(X 
thus: H-O-Cl and H-O— ?As. But, in each case, the 

H-CK 
element which is in combination with hydroxyl has the 
power to combine with oxygen. Hypochlorous acid forms 
the products (HO)CIO, (HO)C10 2 , and (HO)C10 8 , while 
arsenious acid forms arsenic acid (HO) 3 AsO. 

We may consider water as forming the connecting link 
between the oxygen acids and bases. If A stands for any 
acid-forming element, and B for any base-forming element, 
then the general formula of a base is B(OH), and that of 
an oxygen acid A(OH) or O x A(OH), in which O x stands 
for some number of oxygen atoms from one to three or 
four. We should then have these relations; 



160 COLLEGE CHEMISTRY, 

Base<=. Water. Acids. 

I. B'(OH) HOH (O x A)'(OH) 

II. B"(OH) 2 ggg (<W'(OH) a 

HOH 

III. B"'(OH) 3 HOH (O x A)"'(OH), 

HOH 

In these general formulas B" means any bivalent metal, 
and B f " any trivalent metal; and (O x A)" means any group 
of atoms which has the power to hold two hydroxyl groups 
in combination, and is therefore bivalent like (0 2 S), and 
(O x A)' v/ means a trivalent group like (OAs). 

Constitution of Salts. — The view held in regard to the 
constitution of salts is based directly upon those held in 
regard to the constitution of acids and bases. It is believed 
that when an oxygen acid acts upon a base the action takes 
place as represented in the following equation : 



2 N-0-|H + H-0|-K = 2 N-0-K + H-O-H. 

In terms of the theory of solution now generally held 
(see p. 90), an acid is a substance that is dissociated by 
water in such a way as to yield hydrogen ions, while a 
base is dissociated under the same conditions so as to yield 
hydroxyl ions. It is difficult to account for the common 
properties of all acids without assuming that in the solu- 
tions of these acids the common constituent, hydrogen, is 
present in the same condition in all of them. This neces- 
sitates the assumption that when an acid is dissolved in 
water the hydrogen is in some way separated from the rest 
of the molecule, and that from this molecule, therefore, 
two separate parts are formed. These are called ions. 
Acid properties are ascribed to hydrogen ions. In the 
same way the common constituent of bases is hydroxy], 
and this is believed to be separated from the rest of the 
molecule. So that in a solution of a base, no matter what 



BASICITY OF ACIDS. 161 

this base may be, hydroxyl ions are believed to be present, 
and the common properties of the bases are believed to be 
due to the presence of these ions in the solutions of the 
bases. This subject will be discussed more fully farther 
on (see Chapter XXIII). It will then be shown that there 
are good reasons for believing that the extent of dissocia- 
tion of an acid or a base in water solution is dependent 
upon the dilution. The dissociation increases with the 
dilution. 

When a dilute solution of an acid is brought together 
with a dilute solution of a base, the hydrogen ions of the 
acid combine with the hydroxyl ions of the base to form 
water, while the other ions of the acid and of the base 
remain as ions until the solution is concentrated, when 
they unite and finally separate in solid condition as the 
salt. 

Accordingly the action of hydrochloric acid upon potas- 
sium hydroxide is to be represented thus : 



H + C1 + K + 0H = K4-C1 + H 2 0. 
That of nitric acid upon sodium hydroxide thus: 



H + NO s + Na + OH = Na -f N0 8 + H 2 0. 

In the reactions between acids and bases, therefore, the 
chemical act involved is always the union of hydrogen ions 
with hydroxyl ions and the formation of water molecules. 

Basicity of Acids.— In working with acids and bases it 
is noticed that some acids have the power to form but one 
salt with a base like potassium hydroxide, while others 
have the power to form two or more salts with such a base. 
Thus, for example, hydrochloric acid, HC1, and nitric 
acid, HXO3 , can form but one salt with potassium 
hydroxide, and the reactions are represented by the follow- 
ing equations: 



1 62 COLLEGE CHEMISTRY. 

KOH + HC1 = KC1 + H 2 ; 
KOH + HN0 3 = KN0 3 + H 2 0. 

If only half the quantity of base which is required to 
neutralize the acid is added, half the acid remains un- 
changed, and on evaporating the solution the excess of 
acid will pass off. So, also, if only half the quantity of acid 
which is required to neutralize the base is added, half the 
base will remain unchanged. On the other hand, if an 
acid like sulphuric acid is taken, it is found that this has 
the power to form two distinct salts with potassium 
hydroxide, in one of which there is twice as much of the 
metal as in the other. The reactions are represented thus : 

KOH + H 2 S0 4 = KHS0 4 + H 2 0; 
2KOH + H 2 S0 4 = K 2 S0 4 + H 2 0. 

If to a given quantity of sulphuric acid only half the 
quantity of potassium hydroxide which is required to 
neutralize it is added, the first reaction takes place; but if 
the act of neutralization is complete the second reaction 
takes place. An acid of this kind can, further, form one 
salt with two bases, in which one metal is substituted for 
one of the hydrogen atoms of the acid and a second metal 
for the other. 

The different properties of the two kinds of acids 
referred to are ascribed to differences in constitution. In 
the molecule of hydrochloric acid, as in that of nitric acid, 
there is but one atom of hydrogen according to the views 
at present held. If, therefore, the act of neutralization 
takes place in each molecule it is complete, and the salt is 
said to be a neutral or normal salt. In sulphuric acid, 
however, there are two atoms of hydrogen in each mole- 
cule, and either one or both of these may be replaced. If 
only one is replaced, a salt of the general formula MHS0 4 
is obtained. This is still an acid, while also partly a salt. 
It is in fact an acid salt or a salt acid. 

Acids like hydrochloric and nitric acids have not the 
power to form acid salts. They are called monobasic acids. 
While acids like sulphuric acid, which can form two salts 



A CI PITY OF BASES. 163 

with one base, one of which is acid, are called dibasic 
acids. 

Monobasic acids are those which contain but one replace- 
able hydrogen atom in the molecule. Dibasic acids are 
those which contain two replaceable hydrogen atoms in 
the molecule. 

Similarly, there are tribasic acids, like phosphoric acid, 
H 3 P0 4 , arsenic acid, H 3 As0 4 , etc.; tetrabasic acids, like 
pyrophosphoric acid, H 4 P 2 7 ; pentabasic acids, like 
periodic acid, H.I0 6 ; etc., etc. The higher the basicity 
of the acid the greater the variety of salts it can yield. 

Acidity of Bases. — Just as we speak of monobasic, 
dibasic, tribasic acids, etc., so we distinguish between 
bases of different acidity. Thus there are the monacid 
bases, like potassium and sodium hydroxides, KOH and 
NaOH ; diacid bases, like calcium and barium hydroxides, 
Ca(OH) 2 and Ba(OH) 2 ; triacid bases, like aluminium and 
ferric hydroxides, Al(OH) 3 and Fe(OH) 3 ; etc., etc. 

If a monobasic acid acts upon a monacid base, one 
molecule of one forms a salt with one molecule of the 
other, and, in general, no other reaction between the two 
is possible. If a monobasic acid acts upon a diacid base 
two reactions are possible, just as when a monacid base 
acts upon a dibasic acid. Thus, when, for example, 
hydrochloric acid acts upon zinc hydroxide, Zn(OH) 2 , two 
reactions are possible: 

Zn(OH) 2 + HCl =Zn<^ H + H 2 0; 
Zn(OH) 2 + 2HC1 = ZnCl 2 + 2H 2 0. 

The compound ZnCl(OH) is still basic, just as the salt 
KHS0 4 is still acid, and it is called a basic salt. Similarly, 
a triacid base can form three salts with a monobasic acid 
as, for example, in the case of bismuth hydroxide and nitric 
acid, in which three reactions are possible : 

( OH ( NO, 

Bi \ OH + HNO, = Bi^OH + HO; 
OH OH 



1 64 COLLEGE CHEMISTRY. 

{ OH ( N0 3 

Bi \ OH + 2HN0 3 = Bi \ N0 3 + 2H 2 0; 
(OH (OH 

( OH ( N0 8 

Bi 1 OH + 3HN0 3 = Bi \ N0 3 + 3H 2 0. 

( OH ( N0 3 

The salts Bi -J /qt|\ and Bi -J Xjt 3 ^ 2 are basic salts or 

basic nitrates of bismuth, while the salt Bi(N0 3 ) 3 is the 
neutral or normal salt. 

Salts. — From the above it appears that there are three 
classes of salts: (1) Normal salts, which are derived from 
the acids by the substitution of metal atoms for all the 
acid hydrogen atoms; (2) Acid salts, which are derived 
from the acids by the substitution of metal atoms for part 
of the hydrogen; and (3) Basic salts, which are derived 
from the bases by neutralization of part of the basic 
hydroxyl by acids. Normal salts are generally neutral; 
or, if by a neutral substance is meant one which has not 
the power to form salts with acids nor with bases, then the 
expression normal salt is synonymous with neutral salt. 
But some normal salts have what is called an acid reaction, 
and others have an alkaline or basic reaction. Thus a 
normal salt of a weak acid with a strong base as sodium 
carbonate, Na 2 00 3 , has an alkaline reaction. So, also, a 
normal salt of a strong acid with a weak base may have an 
acid reaction, as in the case of aluminium sulphate, 
A1 2 (S0 4 ) 3 . As generally used, the expression neutral salt 
means a salt which exhibits neither an acid nor an alkaline 
reaction, towards certain colored substances such as 
litmus or methyl orange. But a sharp distinction cannot 
be made in this way, as the colored substances used as in- 
dicators do not all act in the same way. The alkaline 
reaction of normal salts of weak acids, and the acid reac- 
tion of normal salts of weak bases is due to the action of 
the water upon these salts. Thus, iu solution, sodium 
carbonate, Na 2 C0 3 , is decomposed to a slight extent by 
the water, and the strong base sodium hydroxide gives the 



ACID PROPERTIES AND OXYGEN. 165 

solution an alkaline reaction. So, also, aluminium sul- 
phate. Al 2 (SOJ 3 , is slightly decomposed by the water, and 
the strong acid, sulphuric aeid, gives the solution an acid 
reaction. A clearer and more satisfactory explanation of 
these phenomena will be given later. 

In naming acid salts various methods are adopted. In 
the ease of a dibasic acid, the only distinction necessary is 
between the acid and the normal salts. The expressions 
acid potassium sulphate and normal potassium sulphate 
mean, of course, the salts which have the formulas KHS0 4 
and K 2 S0 4 , and there is no danger of confusion. We 
may, however, use the names mono-potassium sulphate and 
di'potassium sulphate, or primary and secondary potassium 
sulphates. The last "names are convenient and readily 
convey to the mind the nature of the salts spoken of. Just 
as dibasic acids yield primary and secondary salts, so tri- 
basic acids yield primary, secondary, and tertiary salts. 
For example, phosphoric acid yields three classes of salts: 
primary phosphates, of the general formula MH 2 P0 4 ; 
secondary phosphates, of the general formula M 2 HP0 4 ; and 
tertiary phosphates, of the general formula M 3 P0 4 . The 
phosphates of the first two classes are called, in general, 
acid phosphates. The tertiary phosphate is identical with 
the normal phosphate. In naming basic salts there is no 
difficulty in the simplest cases. Thus, taking the three 
bismuth nitrates the formulas of which are given above, 

the one of the formula Bi -J /qt1\ is called the mono- 
nitrate; that of the formula Bi -j x-rr , the di-nitrate; 

and that of the formula Bi(N0 3 ) 3 , the tri-nitrate, or 
normal nitrate. 

Acid Properties and Oxygen. — Almost all those sub- 
stances which are called acids contain oxygen, as, for 
example, nitric acid, HX0 3 ; sulphuric acid, H 2 S0 4 ; phos- 
phoric acid, H 3 P0 4 ; silicic acid, H 2 Si0 3 ; carbonic acid, 
H 2 C0 3 ; boric acid, H 3 B0 3 ; etc. The presence of oxygen 
in acids was recognized by Lavoisier. As he showed its 



1 66 COLLEGE CHEMISTRY. 

presence in acids to be general, and as lie found that 
several elements and some compounds are converted into 
acids by combination with oxygen, he concluded that this 
element is an essential constituent of all acids, and there- 
fore called it oxygen, a name which, as already stated 
(see p. 32), means the acid-former. According to 
Lavoisier, hydrochloric acid, like other acids, contained 
oxygen, and this view prevailed for many years. As has 
been pointed out under the head of Chlorine, many inves- 
tigations were undertaken with the object of determining 
whether this element does or does not contain oxygen, the 
result being to show that in chlorine, and consequently in 
hydrochloric acid, there is no oxygen. Several acids are 
now known which are like hydrochloric acid in this 
respect, but the latter is the best known example. Similar 
compounds are hydro bromie acid, HBr; hydriodic acid, 
HI; and hydrocyanic acid, HON. The number of these 
acids is, however, quite small, and it is undoubtedly true 
that, of the compounds which we commonly call acids, by 
far the larger number contain oxygen as an essential con- 
stituent. Further, some compounds which are basic can 
be converted into acids by introducing oxygen into them. 
Nomenclature of Acids. — The names of the acids of 
chlorine illustrate some of the principles of nomenclature 
in use in chemistry. The acid of the series which is best 
known is called chloric acid. In naming acids the suffix 
ic is always used in naming the principal member of a 
group of acids containing the same elements. This is seen 
in the names hydrochloric, sulphuric, nitric, phosphoric, 
silicic, carbonic, acetic, etc. If there- are two acids con- 
taining the same elements, that one of the two which 
contains the smaller proportion of oxygen is given a name 
ending in ous. Thus we have the two series: 

Chloric acid HC10 3 Chlorous acid HC10 2 

Sulphuric acid .... H 2 S0 4 Sulphurous acid H 2 SQ 3 

Nitric acid HN0 3 Nitrous acid UNO, 

Phosphoric acid. . . H 3 P0 4 Phosphorous acid. . . . H 3 P0 3 



NOMENCLATURE OF BASES. 167 

For most cases which present themselves this method of 
naming will suffice, but in others the number of acids 
known is larger than two, as, for example, in the series of 
chlorine acids. In such cases recourse is had to prefixes. 
If there is an acid known containing a smaller proportion 
of oxygen than the one whose name ends in ous, it is 
generally designated by means of the prefix hypo, which is 
derived from the Greek vno, signifying under. Thus 
there are the following examples: Hypochlorous acid, 
HC10; hypo-sulphurous acid, H 2 S0 2 ; hyponitrous acid, 
H 2 N 2 2 ; and hypophosphorous acid, H g P0 2 . It will be 
seen on comparing the formulas of these acids with those 
above given that they differ from them in a very simple way. 

In the series of chlorine aeids there is one which contains 
a larger proportion of oxygen than chloric acid. It is 
called perchloric acid, the Latin prefix per signifying here 
very or fully. Similarly there is a perbromic acid and a 
permanganic acid. Other cases arise, but they are of a 
more or less special character, and the compounds are 
given special names according to circumstances. 

Nomenclature of Bases. — As pointed out above, a base 
is a compound of a metal with hydrogen and oxygen. 
The bases are commonly known as hydroxides; and in 
order to distinguish between the hydroxides of the different 
metals, the names of the metals are put before the name 
hydroxide, as in naming the oxides and chlorides. Thus, 
as has been seen, caustic soda, NaOH, is called sodium 
hydroxide, etc. It is necessary in some cases to distinguish 
between two hydroxides of the same metal. This is done 
by using the suffixes ous and ic in the same sense as they 
are used in naming oxides and chlorides. Thus ferric 
hydroxide has the composition Fe(OH) 3 , and ferrous 
hydroxide the composition Fe(OH) 2 ; cuprous hydroxide is 
Cu(OH), and cupric hydroxide Cu(OH) 2 , etc. These 
compounds are sometimes called hydrates, and there are 
some good reasons for using this name, as will be more 
fully shown later. On the other hand, compounds in 
which water as such is regarded as present are called 



168 COLLEGE CHEMISTRY. 

hydrates, and there is danger of confusion if the same 
name is used to designate what are believed to be two 
entirely different glasses of compounds. As examples of 
hydrates we have salts with their water of crystallization, 
chlorine hydrate, Cl 2 + 8H 2 ; hydrochloric acid hydrate, 
HC1 -|- 2H 2 ; etc. While some of the compounds which 
are commonly regarded as hydrates should probably be 
classed with the hydroxides, there seem to be two classes, 
and it is therefore desirable to have two names. 

Nomenclature of Salts. — Theoretically every metal can 
yield a salt with every acid. The salts derived from a 
given acid receive a general name, and this general name 
is qualified in each case by the name of the metal contained 
in the salt. Thus, all the salts derived from nitric acid 
are called nitrates; all the salts derived from chloric acid 
are called chlorates; the salts of sulphuric acid are called 
sulphates; * the salts of phosphoric acid are called phos- 
phates; * etc. So, too, further, the salts of chlorous acid 
are called chlorites; those of nitrous acid, nitrites; those 
of sulphurous acid, sulphites; etc., etc. It will be noticed 
that the final syllable of the name of the salt differs 
according to the name of the acid. If the name of the 
acid ends in ic, the name of the salt derived from it ends 
in ate. If the name of the acid ends in ous, the name of 
the salt ends in ite. To distinguish between the different 
salts of the same acid, the name of the metal contained in 
it is prefixed. Thus, the potassium salt of nitric acid is 
called potassium nitrate, the sodium salt is called sodium 
nitrate; the calcium salt of sulphuric acid is called calcium 
sulphate; the magnesium salt of nitrous acid is magnesium 
nitrite; the calcium salt of hypochlorous acid is calcium 
hypochlorite; etc., etc. If a metal forms two salts with 
the same acid in one of which the valence of the metal is 
lower than in the other, the one in which the valence of 
the metal is lower is designated by means of the suffix ous, 

* Strictly speaking, the salts of sulphuric acid should be called 
sulphurates, and those of phosphoric acid phosphorates, but for the 
sake of euphony these names are shortened to the above forms. 



I 



NOMENCLATURE OF SALTS. 169 

while the one in which the valence of the metal is higher 
is designated by means of the suffix ic. Thus there are 
two series of salts of iron which correspond to the two 
chlorides FeCl 2 and FeCl 3 . In one series the iron appears 
to be bivalent, in the other trivalent. Examples are, 
Fe(N0 3 ) 2 and Fe(N0 3 ) 3 ; FeS0 4 and Fe 2 (S0 4 ) 3 ; etc. 
Those salts in which the iron is bivalent are called ferrous 
salts, as ferrous nitrate, ferrous sulphate, etc. ; and those 
in which it is trivalent are called ferric salts, as ferric 
nitrate, ferric sulphate, etc. Similarly there are two series 
of copper salts known as cuprous and cupric salts; and 
two series of mercury salts known as mercurous and 
mercuric salts. 

If the salts of hydrochloric acid were named in accord- 
ance with the principle just explained, they would be 
called hydrochlorates, and this name is sometimes used for 
complex salts, but in the case of the salts of the metals it 
will be observed that these are identical with the products 
formed by direct combination of the metals with chlorine. 
Thus, hydrochloric acid and zinc act as represented in the 
equation 

Zn + 2HC1 = ZnCl 2 + H 2 ; 
while zinc and chlorine act thus : 

Zn + Cl 2 = ZnCl 2 . 

In each case the same product, Zn01 2 , is formed. But 
these compounds of metals with chlorine are called 
chlorides, as has already been explained. Hence for these 
cases the name hydroclilorate is unnecessary. 

The name hydrate to which reference was made in a 
paragraph above suggests a salt of hydric acid. Potassium 
hydrate signifies the potassium salt of this acid or of water. 
In one sense this is a proper name for the compound. It 
is water in which a part of the hydrogen is replaced by a 
metal, and it is in this respect like a salt. While, how- 
ever, there is an unmistakable analogy between the forma- 
tion of a metallic hydroxide from water and that of a salt 
from an acid, it appears, on the whole, wise not to class 



1)6 



COLLEGE CHEMISTRY. 



water with the acids nor with the bases, but rather to 
regard it as the connecting link between the two classes. 
We shall see later that the similar compounds hydrogen 
sulphide, H 2 S, and hydrogen selenide, H 2 Se, have much 
more marked acid properties than water. When treated 
with metallic hydroxides they form salts of the general 
formulas M 2 S and M 2 Se. 



EXPERIMENTS. 
JSTeutkalization of Acids and Bases; Formation of 

Salts. 
Experiment 77. — Make dilute solutions of nitric, hydrochloric, 
and sulphuric acids (1 part dilute acid, such as is used iu the 
laboratory, to 50 parts water), and 
of caustic soda and caustic potash 
(about 1 gram to 200 cc. of water). 
Measure off a definite volume, say 
20 cc, of each of the acid solu- 
tions. Add a few drops of a solu- 
tion of blue litmus.* Gradually 
add to each of the measured quan- 
tities of acid sufficient dilute caus- 
tic soda to cause the red color 
just to change to blue. As long 
as the solution is red it is acid. 
When it turns blue it is alkaline. 
At the turning-point it is neutral. 
The operation is best carried on 
by means of a burette, which is a 
graduated tube with an opening 
from which small quantities can 
be poured. A convenient shape is 
that represented in Fig. 39. At 
the lower end is a small opening. 
The flow of the liquid from the 
burette is controlled by means of 
a small pinch-cock. It will require 
some practice to enable the student 
to know exactly when the red 

* Methyl orange may be used instead of litmus, and it gives some- 
what sharper results. In alkaline solution this has an orange color. 
Acids change this to rose-red. 




Fig. 39. 



NEUTRALIZATION EXPERIMENTS. *7* 

color disappears and the blue appears, but with practice the 
point can be discerned with great accuracy. Should too much 
alkali be allowed to get into the acid, add a small measured 
quantity of the acid from another burette. Having in one ex- 
periment determined how much of the solution of alkali is re- 
quired to cause the red color to change to blue in operating on a 
given quantity of the acid solution, try the experiment again, 
using a different quantity of the acid solution. If the results of 
several experiments with the same acid and alkali are recorded, 
it will be found that there is a definite ratio between the quanti- 
ties of acid and alkali solution required to neutralize one another. 
If, for example, 15 cc. of the alkali solution are required to neu- 
tralize 20 cc. of the acid solution, 18 cc. of the alkali solution will 
be required to neutralize 24 cc. of the acid solution, 30 cc. to 
neutralize 40 cc, etc. In other words, in order to neutralize a 
given quantity of an acid, a definite quantity of an alkali is nec- 
essary. Perform similar experiments with the other acids. 
Afterwards carefully examine the numerical results. Suppose it 
should require 15 cc. of the caustic-soda solution or 12 cc. of the 
caustic-potash solution to neutralize 20 cc. of the hydrochloric- 
acid solution. Compare the quantities of these alkali solutions 
necessary to neutralize equal quantities of the other acids. What 
conclusion is justified with reference to the act of neutralization ? 



Study of the Products formed. 

Experiment 78. — Dissolve about 10 grams caustic soda in 
100 cc. water. Add hydrochloric acid slowly, examining the 
solution from time to time by means of a piece of paper colored 
blue with litmus. As long as the solution is alkaline it will 
cause no change in the color of the paper. The instant the point 
of neutralization is passed, the solution changes the color of the 
paper to red ; when exactly neutral, it will neither change the blue 
to red, nor, if the color is changed to red by means of another 
acid, will it change it back again. When this point is reached, 
evaporate to complete dryness on the water-bath, and see what 
is left. Taste the substance. Has it an acid taste ? Does it 
suggest any familiar substance ? If it is sodium chloride, how 
ought it to conduct itself when treated with sulphuric acid? 
Does it conduct itself in this way ? Satisfactory evidence can be 
given that the substance is sodium chloride. It is not an acid 
nor an alkali. It is neutral. 



172 COLLEGE CHEMISTRY. 

Experiment 79. — Perform a similar experiment, using dilute 
nitric acid and caustic soda. What evidence have you that the 
product in this case is different from caustic soda ? 

Experiment 80. — Perform similar experiments with dilute sul- 
phuric acid and caustic soda ; with sulphuric acid and caustic 
potash ; with nitric acid and caustic potash ; with hydrochloric 
acid and caustic potash. Dry and examine the product carefully 
in each case ; and keep for future study what is not used in 
these experiments. 



CHAPTER XL 

NATURAL CLASSIFICATION OF THE ELEMENTS.— 
THE PERIODIC LAW.* 

Historical. — It has long been known that simple rela- 
tions exist between the atomic weights of some elements 
which resemble one another closely. Thus chlorine, 
bromine, and iodine are very similar elements. Their 
atomic weights are 35.45, 79.96, and 126.85 respectively. 
It will be seen that the atomic weight of bromine, 79.96, 
is approximately the mean of those of chlorine and iodine. 
We have 

35.45 + 126.85 



2 



= 81.15. 



A similar group is that of sulphur, selenium, and tel- 
lurium, which resemble one another as closely as chlorine, 
bromine, and iodine do. The atomic weights are S = 
32.06, Se = 79.1, and Te = 127. We have here 

32.06 + 127 =m3> 

Other groups are those of phosphorus, 31, vanadium, 
51.2, and arsenic, 75: 

lithium 7.03, sodium, 23.05, and potassium, 39.15. 

7 - 03 + 39 - 15 = 23.09. 

2 

Arrangement of the Elements. — MendeleefT (1869) and 
Lothar Meyer (1870) have proposed several arrangements 

* A large table of the Natural System of the Elements should be 
hung up in a conspicuous place in the laboratory. 

173 



174 



COLLEGE CHEMISTRY. 





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THE PERIODIC LAW. 
MENDELEEFF'S TABLE II. 



175 









I. 


II. 


III. 


IV. 


V. 


VI. 


R 2 


I. 




Li= 7 


K 


39 


Rb 85 


Cs 


133 


- - 


RO 


II. 




Be= 9 


Ca 


40 


Sr 87 


Ba 


137 


- - 





R 2 3 


III. 




B = 11 


Sc 


44 


Y 89 


La 


138 


Yb 173 





R0 2 


IV. 


(H 4 C) 


C = 12 


Ti 


48 


Zr 90 


Ce 


142 


- - 


Th231 


R 2 5 


V. 


(H 3 N) 


N '= 14 


V 


51 


Cb 94 


Di 


146 


Ta 182 





R0 3 


VI. 


(H 2 0) 


= 16 


Cr 


52 


Mo 96 


- 


- 


W 184 


U 240 


R 2 7 


VII. 


(HF) 


F = 19 


Mn 


55 












R0 4 








Fe 


56 


Ru 103 


- 


- 


Os 192? 







VIII. 






Co 






- 


- 















Ni 


59 


Pd 106 


- 


- 


Pt 195 


- - 


R a O 

RO 


I. 

II. 


H = 1 


Na = 23 


Cu 
Zn 


63 
65 


Ag 108 
Cd 112 


— 


— 


Au 196 
Hg 200 


:: 




Mg 24 


R 2 3 


III. 




Al 27 


Ga 


69 


In 113 


- 


- 


Tl 204 


— 


R0 3 


IV. 


(H 4 R) 


Si 28 


Ge 


72 


Sn 118 


- 


- 


Pb 206 


— 


R 2 5 


V. 


(H 3 R) 


P 31 


As 


75 


Sb 120 


- 


- 


Bi 209 


- - 


R0 3 


VI. 


(H 2 R) 


S 32 


Se 


79 


Te 125? 










R 2 7 


VII. 


(HR) 


CI 35.5 


Br 


80 


I 127 











for the purpose of making clear the connection between 
the properties and atomic weights of the elements. Those 
which have proved most useful will first be given, and then 
the connection between the atomic weights and properties 
will be discussed briefly. The different arrangements are 
to be regarded only as different ways of expressing the 
same law, and no one of them is perfect. The investiga- 
tion of the relations between the atomic weights and the 
properties of the elements has not yet been pushed far 
enough to justify a final opinion as to the character of the 
relations, but it has nevertheless reached a stage in which 
we are justified in stating that these relations are general 
and deep-seated. 

In the above tables the approximate atomic weights are 
used instead of those which have been determined and 
calculated with the greatest care. For most purposes in 



176 COLLEGE CHEMISTRY. 

the laboratory the approximate figures answer well enough, 
and they are commonly used. In the following table of 
Lothar Meyer the refined atomic weights are used. The 
difference between the two sets of figures is in most cases 
slight. 

In MendeleefFs Table I the elements are arranged in 
horizontal lines, beginning with lithium. When the 
eighth element, sodium, in the order of the increasing 
atomic weights is reached it is found that it is very much 
like lithium. If this is placed below lithium, and the 
next six elements in the same horizontal line, when the 
fifteenth element is reached, it is found like the eighth to 
be similar to lithium. Up to and including manganese 
there are twenty-one* elements excluding hydrogen. 
These fall then naturally into three series of seven mem- 
bers each, and placing these horizontally, those elements 
which fall in the same perpendicular lines have the same 
general character. This is seen most strikingly in 
Group I, in which lithium, sodium, and potassium fall, 
and in Group V, in which nitrogen, phosphorus, and 
vanadium fall; but there is no difficulty in recognizing the 
similarity in the other groups. The three elements fol- 
lowing manganese, viz., iron, nickel, and cobalt, are very 
much alike, and they certainly do not belong in Groups 
I, II, and III, while the next element, copper, has some 
properties which ally it to the members of Group I. The 
next six elements fall in Groups II to VII, and are evi- 
dently in place, and the six following fall in Groups I to 
VI, and are also in their proyjer places, as far as their 
properties are concerned. After molybdenum in the sixth 
series comes a blank which means that there is no element 
to fill that place, but that probably there is one undis- 
covered which has the atomic weight approximately 100, 
and has properties similar to those of manganese. Then 

* Since tliis table was prepared by Mendeleeff three additional 
elements Lave been discovered that have atomic weights lower than 
that of manganese. These are helium (at. wt. 4), neon (at. wt. 20), 
and argon (at. wt. 40), 



THE PERIODIC LAW. 



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173 COLLEGE CHEMISTRY. 

follow three elements which resemble one another as 
closely as iron, nickel, and cobalt do. These do not 
belong in Groups I, II, and III, but form a small inde- 
pendent group. These two groups of three elements 
occur at the end of the fourth and sixth series respectively. 
We should therefore expect to find a similar group at the 
end of the eighth series. No such group is known, how- 
ever, though at the end of the tenth series, where we 
should look for the next similar small group, there are 
the three elements iridium, platinum, and osmium. The 
elements of series 2 beginning with lithium and ending 
with fluorine differ in some respects quite markedly from 
all the other elements, as will be seen when they are taken 
up. Beginning with sodium, it will be seen that there 
are two series of seven elements and a short series of three ; 
then again two series of seven and a series of three ; and, 
although the following series are imperfect, it is not diffi- 
cult to recognize that the same general arrangement of 
the elements holds good to the end. A series of seven 
elements is called a short peri oil; while two short periods 
with the accompanying three similar elements constitute 
what is called a long period. 

In Mendeleeff's Table II the long periods are arranged 
in perpendicular lines, each long period beginning and 
ending with a short period and having a side group of 
three elements in the middle. Thus in the column 
beginning with potassium there is, first, the short period 
potassium to manganese, then the side group iron, cobalt, 
nickel, and then the short period copper to bromine. In 
this table similar elements occur in the same horizontal 
lines. Thus in one line there are lithium, potassium, 
rubidium, and cesium; in another sulphur, selenium, and 
tellurium: and in another chlorine, bromine, and iodine. 

The symbols at the top of each column in Table I have 
reference to the general formulas of the compounds which 
the elements in each group form with oxygen and with 
hydrogen. Beginning with Group I, the general formula 
of the oxygen compounds of the members of this group is 



THE PERIODIC LAW. 179 

E..0, in which R represents any element of that group; 
the general formula of the oxygen compounds of the 
members of Group II is RO; and so on. It will be 
observed that the oxygen compounds grow more and more 
complex from Group I to Group VII. Writing RO, R0 2 , 
and RO., with doubled formulas, thus: R 2 2 , R,0 4 , and 
R.,0 6 , the series of oxygen compounds is represented as 
below : 

R 2 0, R 2 2 , R 2 O s , R 2 Q 4 , R 2 5 , R 2 6 , R,O r 

As regards the general formula of the oxygen com- 
pounds of the members of Group VIII, it must be said 
that it does not in general correspond to the composition 
R0 4 . Osmium and ruthenium do, however, form the 
oxides Os0 4 and Ru0 4 . 

There is also regularity in the composition of the 
hydrogen compounds. Beginning with Group VII, those 
members which combine with hydrogen form compounds 
of the general formula RH, as, for example, C1H, hydro- 
chloric acid; FH, hydrofluoric acid; etc. Those members 
of Group VI which combine with hydrogen form com 
pounds of the general formula RH 2 , as, for example, 
water, H 2 0, and lrydrogen sulphide, H 2 S. The maximum 
power of combining with hydrogen is met with in Group 
IV, in which occur the elements carbon and silicon. 
These form the hydrogen compounds CH 4 and SiH 4 . The 
members of Groups I, II, and III do not readily form 
compounds with hydrogen. A few are known, but they 
are quite unstable. 

The hydroxides vary in composition from the simple 
form R(OH) to R(OH)_. While in the first four groups 
well-marked examples of the hydroxides R(OH), R(OH) 2 , 
R(OH) 3 and R(OH) 4 are found, in the fifth group tlie 
hydroxides have not the general formula R(OH)., though 
several of them have the formula OR(OH) 3 , as phosphoric, 
arsenic, and antimonic acids, which are respectively 
OP(OH) 3 , OAs(OH) 3 , and OSb(OH) 3 . These may be 



180 COLLEGE CHEMISTRY. 

regarded as derived from the hydroxides of the general 
formula R(OH) 5 by loss of water: 

B(QH), = OBCPH), + H 1 0. 

Hydroxyl derivatives of the members of Group VI corre- 
sponding to the general formula E(OH) 6 are known, as, for 
example, the so-called hydrate of sulphuric acid, S(OH) 6 . 
The maximum hydroxides of Group VII should have the 
general formula R(OH) 7 , but those known do not corre- 
spond to this. The nearest approach to it is found in 
crystallized periodic acid, H 5 I0 6 , which may be regarded 
as derived from the hydroxide I(OH) 7 by loss of one mole- 
cule of water, thus : 

I(OH), = OI(OH), + H 2 0. 

The arrangement of Lothar Meyer is a continuous one. 
The elements are arranged on a spiral beginning with 
lithium and ending with uranium. The divisions are 
such that when the two ends of the table are brought 
together on a cylinder, the line ending with fluorine will 
join that beginning with sodium; that ending with nickel 
will join that beginning with copper; and so on. In other 
respects the arrangement is much like that in MendeleefFs 
Table I. What Mendeleeff calls a Group, Lothar Meyer 
calls a Natural Family, while those elements which fall 
in the same horizontal line are said to form a series. 
Now, each natural family falls into two groups indicated 
by the letters A and B placed above. Thus the first 
natural family falls into Group A, consisting of lithium, 
sodium, potassium, rubidium, and caesium, and Group B, 
consisting of copper, silver, and gold. Family II falls into 
Group A, consisting of glucinum, magnesium, calcium, 
strontium, barium, and perhaps erbium, and Group B, 
consisting of zinc, cadmium, and mercury; etc. The 
members of each group in a family resemble one another 
much more closely than they resemble the members of the 
other group. 



THE PERIODIC LAW. 181 

The figures at the bottom of the table of Lothar Meyer 
refer to the valence of the elements in each group. Judg- 
ing by the composition of the oxides and hydroxides, the 
valence increases from 1 to 8 from Families I to VIII. 
But the valence of the elements in each family varies 
according to conditions. Thus the valence of the elements 
of Family IV is generally 4, as shown in the compounds 
CH 4 , CC1 4 , C0 2 , SiH 4 , SiCl 4 , Si0 2 , etc.; but they may 
also appear as bivalent elements, as seen in the compound 
CO. So, too, while the elements of Family V are quin- 
quivalent, as in PCL, NH 4 C1, etc., they may also be tri- 
valent, as in PC1 3 , NH 3 , etc. 

What we call valence does not then appear to be an 
unchangeable property of the elements, but a property 
which may change according to conditions. It appears 
further that a given element may have one valence towards 
one element and another valence towards another element. 
This is most strikingly seen on comparing the formulas of 
the hydrogen compounds of the elements of Families V, 
VI, and VII with those of their oxygen compounds. The 
members of Family VII combine with hydrogen in only 
one proportion, and that is the simplest possible. Towards 
hydrogen these elements are univalent, and their valence 
towards hydrogen is constant. On the other hand, they 
combine with oxygen and with hydroxyl in several propor- 
tions, and judging by the composition of these compounds, 
the valence towards oxygen varies from 1 to 7. The 
members of Family VI are bivalent towards hydrogen, and 
their hydrogen valence is constant ; but they combine with 
oxygen and hydroxyl in several proportions, and the com- 
position of the compounds indicates that their valence 
towards oxygen varies from 2 to 6. The hydrogen valence 
of the members of Family V is 3, while the oxygen valence 
varies from 1 to 5. Finally, the hydrogen valence of the 
members of Family IV is 4, while the oxygen valence 
varies from 2 to 4. As regards the hydrogen valence of 
the members of Families I to III, but little is known. 
These elements do not generally combine with hydrogen, 



1 82 COLLEGE CHEMISTRY. 

though some of them do. Towards oxygen their valence 
is fairly constant, though some variations are observed as 
in the case of copper and mercury. 

Judging then by the composition of the compounds, we 
are justified in making a distinction between the hydrogen- 
valence of some elements and their oxygen-valence. While 
the former is constant, the latter is subject to variations. 
In those cases in which there is a marked difference 
between the liydrog en-valence of an element and its maxi- 
mum oxygen- valence, the maximum valence towards chlo- 
rine is greater than the hydrogen-valence and less than the 
maximum oxygen-valence. This is shown in the case of 
sulphur; the formulas of its hydrogen compound and of 
its highest compounds with chlorine and oxygen being 
respectively 

SH 2 , SC1 4 , S0 3 . 

From this it appears that the maximum valence of sulphur 
towards hydrogen is 2, towards chlorine 4, and towards 
oxygen 6. On the other hand the valence of sulphur 
towards fluorine is 6 as is shown by the stable compound 
sulphur hexafluoride, SF 6 . 

Connection between the Position of the Elements in the 
Natural System and their Chemical Properties. — The 
changes in composition of the oxygen and hydrogen com- 
pounds and of the hydroxides from Family I to VII have 
been referred to. Another fact of great importance is that 
the elements of Group I are the most strongly marked 
base-forming elements, while those of Group VII are the 
most strongly marked acid-forming elements. Passing in 
either direction, the character of the elements becomes less 
pronounced, until in the middle (Group IV), elements 
which form neither strongly-marked acids nor ^ strongly- 
marked bases are found. Thus, beginning with sodium, 
this element forms a strong base, magnesium forms a 
weaker base, the hydroxide of aluminium is a still weaker 
base. Beginning, on the other hand, with chlorine at the 
other end of the same series, its hydrogen compound is a 



THE PERIODIC LAW. 183 

strongly-marked acid; that of sulphur is an acid, but less 
marked in character than hydrochloric acid; that of phos- 
phorus has no acid properties, nor has that of silicon. 
The hydroxides of these four elements have acid properties. 
Each one, however, forms several acids, and it is difficult 
to compare them, as some of those of chlorine are strongly 
marked and others are not, as we have seen. 

Some very interesting variations in properties are also 
noticed in passing from one end of a group of a natural 
family to the other. Thus in Group B, Family VII, the 
activity of the elements grows less from fluorine to iodine, 
or, as we commonly say, fluorine is the strongest element 
in the group, and then follow, in order, chlorine, bromine, 
and iodine. 

The remarkable relations above referred to are summed 
up in the periodic law: 

The properties of an element are periodic functions of the 
atomic iceiglit. 

It appears that if an element has a certain atomic 
weight it must have certain properties, and that if the 
atomic weight is known the properties can be stated, just 
as, if the properties are known, the atomic weight can be 
approximately stated. When the law was first stated, 
Mendeleefl predicted the discovery of certain elements to 
fill some of the vacant places in the table. At that time 
the elements gallium, Ga, scandium, Sc, and germanium, 
Ge, were not known. Not only was their discovery pre- 
dicted, but their properties were clearly stated years before 
they were- brought to light. Within the last few years 
these three elements have been discovered, and a remark- 
able agreement is observed between their properties as 
determined by observation and as foretold by Mendeleeff 
by the aid of the periodic law. 

The relations between the atomic weights and properties 
will appear more and more clearly as our knowledge of 
the elements advances. The natural arrangement of the 
elements suggested by the periodic law is adopted in this 
book. The elements hydrogen, oxygen, and chlorine were 



1$4 COLLEGE CHEMISTRY. 

studied at the outset in order to illustrate the methods of 
studying chemical problems, and as examples of chemical 
elements in general. It is, however, now time to take up 
the elements systematically, and to learn what may be 
necessary in regard to them in order to get as clear a notion 
as possible of the facts and principles of the science of 
chemistry. 

Plan to be followed. — The most systematic method of 
procedure in studying the elements would be to begin with 
Family I, Group A (see Lothar Meyer's Table, p. 178), 
then to take up Group B of the same family; and so on in 
order, ending with Family VIII. It seems better, how- 
ever, to begin with Family VII; to follow with Families 
VI, V, and IV; and then to take up in order Families I, 
II, III, and VIII. The main reason for this is that it is 
impossible to study most of the members of Families 
I, II, III, and VIII without a knowledge of several of the 
elements of Families VII, VI, V, and IV, while these last 
families can be studied with only slight reference to the 
others. It is proposed then to begin with Group B, 
Family VII, the members of which are very much like 
chlorine. The only member of Group A of this family is 
manganese. While manganese resembles the members of 
the chlorine group in some respects, it has other properties 
which ally it to the so-called base-forming elements. So 
also the members of Group A, Family VI, are like the 
members of the oxygen or sulphur group, but they are also 
allied to the base-forming elements. A similar difference 
is observed between the members of Groups A and B, 
Family V. 

"While the plan above sketched takes into consideration 
the greater number of analogies of the elements, there are 
other analogies that are not brought out. Thus, as will 
be seen in due time, the elements aluminium, chromium, 
manganese, and iron are analogous in some respects, but 
by following the plan sketched, they will be taken up in 
different groups. This appears to be justified, however, 
when we consider the entire conduct of these elements, 



THE PERIODIC LAW. 185 

and do not confine ourselves to a study of a few reactions 
which, being useful for analytical purposes, have been 
studied more carefully than others which from a scientific 
point of view are perhaps just as important. 

But, making all concessions, it must be acknowledged 
that the periodic law gives only an imperfect view of the 
relations between the elements. That there is some close 
connection between the properties of an element and its 
atomic weight appears to be evident, and until we have 
some better guide to the classification of the elements than 
the periodic law, it seems wise to follow this. 



CHAPTER XII. 

THE ELEMENTS OF FAMILY VII, GROUP B: 
FLUORINE.— CHLORINE.— BROMINE.— IODINE. 

General. — The elements of this group are commonly 
called the halogens. The best known member of the group 
is chlorine. Although fluorine is in general like the other 
members of the group, it differs from them in some 
respects. While chlorine, bromine, and iodine accompany 
one another in nature, fluorine compounds are not gen- 
erally found in company with compounds of the other 
elements of the family. In those cases in which chlorine, 
bromine, and iodine are found together, chlorine is gen- 
erally present in largest quantity, and iodine in smallest 
quantity. Fluorine and chlorine are gases under ordinary 
conditions, while bromine is a liquid and iodine is a solid. 
Fluorine, bromine, and iodine form w r ith hydrogen the 
compounds hydrofluoric acid, HF, hydrobromic acid, 
HBr, and hydriodic acid, HI, which are analogous to 
hydrochloric acid. All these compounds are gases which 
have marked acid properties. With oxygen, flourine does 
not combine, whereas chlorine, bromine, and iodine com- 
bine with it in a number of proportions, as has already 
been seen in the case of chlorine. Among themselves these 
elements also form some compounds: thus bromine and 
chlorine form the compound BrCl; iodine forms the com- 
pounds IC1, IC1 3 , IBr, and IF 5 . It appears from this 
that the valence of iodine towards bromine is 1, towards 
chlorine 3, and towards fluorine 5. 

Towards base-forming members the elements of this 
group are univalent, as shown in such compounds as NaCl, 
KBr, CaCl 2 , KI, etc. 

186 



BROMINE. 187 

Bromine, Br (At. Wt. 79.96). 

Occurrence. — Bromine occurs in nature in company with 
chlorine. Chlorine, as has been stated, occurs mostly in 
combination with sodium, as sodium chloride, or common 
salt. In several of the great salt-beds bromine occurs in 
the form of sodium bromide, NaBr, and in some places it 
occurs as magnesium bromide, MgBr 2 . The chief source 
01 bromine is the mother-liquors from the salt-works. 
When a solution containing a large quantity of sodium 
chloride and a small quantity of the bromide is evaporated, 
the chloride is first deposited, and from the mother-liquors 
the bromide mixed with chloride is deposited. The great 
beds at Stassfurt are particularly rich in bromides, and a 
great deal of bromine is made from the salts which occur 
in this locality. 

Preparation. — Bromine can be prepared from the 
bromides in the same way that chlorine is made from the 
chlorides. The reaction is represented by the equation 

2NaBr + Mn0 2 + 2H 2 S0 4 = 

Na 2 S0 4 + MnS0 4 + 2H 2 + Br 2 . 

Another method for the preparation of bromine depends 
upon the fact that chlorine has the power to set bromine 
free from its compounds. If, therefore, a solution con- 
taining a bromide is treated with manganese dioxide and 
hydrochloric acid, the chlorine which is formed from the 
hydrochloric acid will act upon the bromide and bromine 
will be given off. This method is used at Stassfurt. 

Properties. — Bromine is a heavy, dark-red liquid at 
ordinary temperatures. If exposed to the air it is con- 
verted into a vapor of a brownish-red color. It boils at 
58-58.6°, and at — 7.3° it is solid. It has an extremely 
disagreeable odor, to which fact it owes its name (from 
fipGOfj.os, a stench). From carbon disulphide at — 90° it 
crystallizes in fine dark-red needles. 

Its properties are similar to those of chlorine. It acts 



i88 COLLEGE CHEMISTRY. 

violently upon organic substances; attacking the skin, and 
the membranes lining the passages of the throat and 
lungs. Wounds caused by the liquid coming in contact 
with the skin are painful and serious, and it must there- 
fore be handled with great care. 

Like chlorine, bromine is dissolved by water, one part 
dissolving in 33.3 parts at 15°. The solution, which has 
a reddish color and the odor of bromine, is called bromine 
water. At a low temperature bromine forms with water a 
compound in every way analogous to chlorine hydrate, 
viz., bromine hydrate, Br 2 -j- 10H 2 O. This decomposes 
when left in contact with the air at ordinary temperatures. 

Chemical Conduct of Bromine. — Bromine acts chem- 
ically like chlorine. It was pointed out that chlorine acts 
in three different ways: (1) By direct addition; (2) by 
substitution; and (3) by liberating oxygen from water, as 
in bleaching and other oxidizing processes. Bromine is 
capable of acting in all three ways. 

Uses of Bromine. — Bromine and its compounds are used 
in photography, medicine, and to some extent in the 
manufacture of coal-tar colors. It is manufactured in 
large quantity, and a large proportion of it is manufactured 
in the United States. According to the official report the 
production of bromine in the United States in the year 
1896 amounted to over 500,000 pounds. 

Hydrobromic Acid, HBr. — The only compound which 
bromine forms with hydrogen alone is hydrobromic acid. 
This is in all respects very much like hydrochloric acid. 
It is set free from bromides by the action of sulphuric 
acid, but owing to its instability it acts upon the sulphuric 
acid, causing decomposition. Hydrochloric acid does not 
act upon sulphuric acid at all. Hydrobromic acid acts 
according to the following equation : 

2HBr + H 2 S0 4 = 2H 2 + S0 2 + Br 2 . 

The action consists in the decomposition of the hydro- 
bromic acid into bromine and hydrogen, and the subse- 



BROMINE AND ITS COMPOUNDS. 189 

quent action of the nascent hydrogen upon the sulphuric 
acid thus: 

2HBr = 211 + Br.,; and 

H 2 S0 4 -f- 2H = 2H 2 4- S0 2 . 

The hydrobromic acid acts here, then, as a reducing 
agent, and the sulphuric acid as an oxidizing agent. It is 
plain that hydrobromic acid cannot be made in pure con- 
dition by the action of sulphuric acid upon a bromide. 
Some of the hydrobromic acid, to be sure, escapes the 
action of the sulphuric acid, but at best it is always mixed 
with the compound S0 2 , or sulphur dioxide, which is a 
gas, and with bromine. 

The method most commonly adopted in the laboratory 
for the preparation of lrydrobromic acid consists in treat- 
ing phosphorus with bromine and water. In all prob- 
ability the bromine acts first upon the phosphorus, forming 
the product PBr 3 or PBr 5 according to the proportions of 
the substances used. Both these substances are decom- 
posed by water, the first forming phosphorous acid and 
hydrobromic acid, according to this equation: 

( Br HHO 
P \ Br + HHO = P0 3 H 3 + 3HBr, 
( Br HHO 

or PBr 3 + 3H 2 = P0 3 H 3 -f 3HBr; 

the second forming phosphoric acid and hydrobromic acid: 
PBr 5 + 4H 2 = P0 4 H, + 5HBr. 

The gas thus formed can be freed from bromine by 
passing it through a tube containing phosphorus. 

Properties. — Hydrobromic acid is a colorless gas that 
forms fumes in contact with the air in consequence of its 
power of combining with water. It dissolves in water in 
large proportion. The solution conducts itself much like 
hydrochloric acid. 

Compounds of Bromine with Hydrogen and Oxygen. — 
With hydrogen and oxygen bromine forms compounds 



190 COLLEGE CHEMISTRY. 

which closely resemble those which chlorine forms with the 
same elements. They are: Hypobromons acid, HBrO; 
bromic acid, HBr0 3 ; and perhaps perbromic acid, HBr0 4 . 

Iodine, I (At. Wt. 126.85). 

Occurrence. — Iodine, as has already been stated, occurs 
in company with chlorine and bromine in nature, but in 
smaller quantity than these. The relative quantity in 
sea-water is extremely small. The sea-plants, however, 
assimilate it, and the ashes of these plants contain a con- 
siderable quantity of compounds of iodine. It also occurs 
in small quantity in the great beds of soda saltpetre, or 
sodium nitrate, which are found in Chili, South America. 
It occurs in small quantity in combination with silver, and 
also in combination with lead and with mercury. 

Preparation. — The method of obtaining iodine from its 
salts is like that used in making chlorine and bromine from 
the chlorides and bromides. It consists in treating the 
iodides with sulphuric acid and manganese dioxide : 

2KI + Mn0 2 + 2H 2 S0 4 = K 2 S0 4 + MnS0 4 + 2H 2 + I 2 . 

The iodine, although solid at the ordinary temperature, 
is easily volatilized, and if the mixture mentioned is 
heated, iodine vapor passes over and may be condensed in 
appropriately arranged vessels. 

On the large scale iodine is obtained mostly from sea- 
weed. On the coasts of Scotland, Ireland, and France the 
seaweed which is thrown up by storms is gathered, dried, 
and burned. The organic portions are thus destroyed, 
and the mineral or earthy portions are left behind as ashes. 
This incombustible residue is called help. It contains a 
small percentage of potassium iodide, from . 5 to 2 per cent 
according to the seaweed used. The dried weed was 
formerly burned in cavities dug in the earth, but of late 
years the process has in some places been much improved, 
and the yield in kelp increased. 



IODINE. I9 1 

In Scotland the iodine is liberated by means of sulphuric 
acid and manganese dioxide. In France, however, this is 
effected by passing chlorine into the solution containing 
the iodide. If too little chlorine is used all the iodine is 
not separated; if too much, a compound of iodine and 
chlorine is formed, or an iodate, in consequence of the 
oxidizing action of the chlorine on the iodine. 

The iodine which occurs in Chili saltpetre, NaN0 3 , is 
in the form of sodium iodate, NaI0 3 , and iodide, Nal, and 
to some extent as magnesium iodide, Mgl 2 . Most of the 
iodine now in the market is made from this material, and 
the competition created in this way has led to a careful 
study of the process for obtaining iodine from kelp. Sea- 
weed is now collected from certain parts of the ocean 
where it grows in large quantity, vessels being sent out for 
the purpose. 

Properties. — Iodine is a grayish-black crystallized solid. 
At ordinary temperatures it js volatile. According to the 
most reliable determinations it melts at 113-115°, and boils 
at 250°. The vapor has a violet color when mixed with 
air. When in pure condition it is intensely blue. At 
temperatures considerably above the boiling-point the 
specific gravity of iodine vapor is such as to show that its 
molecular weight is approximately 254, or twice the atomic 
weight. As the temperature is raised, however, the specific 
gravity is lowered, until, finally, at a very high tempera- 
ture, it becomes about half what it is at lower tempera- 
tures. This is accounted for by supposing that at the 
lower temperatures the molecules of iodine consist of two 
atoms each, while as the temperature is raised these mole- 
cules are gradually broken down, so that at the tempera- 
ture at which the lowest specific gravity is reached the 
iodine vapor consists of free atoms, or the atoms and 
molecules are then identical, and the specific gravity is 
therefore only half what it is when the molecules consist 
of two atoms. 

Iodine has a characteristic strong taste. It acts upon 
the mucous membranes, but much less energetically than 



192 COLLEGE CHEMISTRY. 

chlorine or bromine. It colors the skin yellowish-brown, 
and acts as an absorbent, causing the reduction of some 
kinds of swellings. 

It dissolves slightly in water, easily in alcohol, and easily 
in a water solution of potassium iodide. The solution in 
alcohol is known as tincture of iodine. It dissolves also 
in carbon disulphide, CS 2 , and in chloroform forming 
solutions which have a beautiful deep-violet color. 

In general, iodine conducts itself chemically like bro- 
mine and chlorine, only it acts in almost all reactions less 
energetically than the other two elements. It combines 
directly with a number of elements, as with hydrogen, 
sulphur, phosphorus, iron, mercury, etc. In presence of 
water it acts as an oxidizer just as chlorine and bromine 
do, but less energetically. Thus it oxidizes sulphurous 
acid, H 2 S0 3 , to sulphuric acid, H 2 S0 4 : 

H 2 S0 3 + I 2 + H 2 = H 2 S0 4 + 2HL 

As a substituting agent it does not act as readily as 
chlorine and bromine, though, iodine substitution-products 
are made in large quantities, particularly in connection 
with the manufacture of dye-stuffs. 

Iodine is used extensively in the dye-stuff industry, in 
photography, and in medicine. One factory in Scotland 
makes on an average 60 tons of iodine a year. 

Hydriodic Acid, HI. — Hydriodic acid cannot be made 
pure by treating an iodide with sulphuric acid. The 
hydrogen of the hydriodic acid acts upon the sulphuric 
acid very readily, and according to the conditions the fol- 
lowing reactions may take place : 

H 2 S0 4 + SHI - 4H 2 + SH 2 + 4I 2 ; 
H 2 S0 4 + 6HI = 4H 2 + S + 3I 2 ; 
H 2 S0 4 -f 2HI = 2H 2 + S0 2 + I 2 . 

On treating potassium iodide with sulphuric acid, there- 
fore, there may be formed, in addition to hydriodic acid 
and free iodine, sulphur dioxide, sulphur and hydrogen 
sulphide. . 



HYDRIODIC ACID— IODIC ACID. 193 

The method adopted for the preparation of hydriodic 
acid is like that used for the preparation of hydrobromic 
acid. 

Properties. — Hydriodic acid is a colorless transparent 
gas like hydrochloric and hydrobromic acids. It also like 
these dissolves in water in large quantity, and when 
brought in contact with the air it forms dense white 
fumes. When boiled the water solution conducts itself like 
those of hydrochloric and hydrobromic acids. 

When heated, the gas begins to decompose at 180°, and 
at higher temperatures the decomposition takes place 
rapidly. The products are simply hydrogen and iodine. 
In consequence of the ease with which hydriodic acid 
breaks down, yielding free hydrogen, it is an excellent 
reducing agent, and it is frequently used in the laboratory 
for the purpose of extracting oxygen from substances. Its 
action upon sulphuric acid has already been spoken of. 
The reason why it acts so well is that the hydrogen is 
separated from the iodine with little expenditure of energy, 
and the hydrogen thus separated is in the nascent state, 
or, as is believed, in the atomic state. 

Iodic Acid, HI0 3 . — This compound is strictly analogous 
to chloric and bromic acids, but differs from them in being 
much more stable. 

When iodine is dissolved in an alkali the reaction which 
takes place is the same as that which takes place with 
chlorine and bromine under like circumstances. A mix- 
ture of the iodide and iodate is formed : 

6KOH + 31, = 5KI +- KI0 3 + 3H 2 0. 

Iodic acid is a crystallized solid, which -when heated to 
170° loses water and is converted into iodine pentoxide, 

2HI0 3 =I 2 5 +H 2 0. 

Its salts have the general formula MI0 3 , though it also 
forms salts MH(I0 3 ) 2 and MH 2 (I0 3 ) 3 . It gives up its 



194 COLLEGE CHEMISTRY. 

oxygen readily and is therefore a good oxidizing agent, just 
as hydriodic acid is a good reducing agent. 

Iodine Pentoxide or Iodic Anhydride, I 2 5 . — This com- 
pound is formed, as was stated in the last paragraph, by 
heating iodic acid to 170°. It is a white solid which is 
easily soluble in water, forming iodic acid. It is decom- 
posed when heated to 300°. It will be observed, therefore, 
that this compound of iodine and oxygen is much more 
stable than any of the compounds of chlorine or bromine 
and oxygen; and it is interesting to note that as the 
stability of the oxygen compounds increases, that of the 
hydrogen compounds decreases. 

Anhydrides, or Acidic Oxides. — An oxide which, like 
iodine pentoxide, forms an acid when dissolved in water, 
or which forms salts by treatment with basic hydroxides, 
is called an anhydride or acidic oxide. The oxides of the 
base-forming elements form bases when dissolved in water, 
and they are, therefore, called basic oxides. As examples 
of acidic oxides or anhydrides, there may be mentioned 
besides iodic anhydride, sulphuric anhydride, S0 3 ; sul- 
phurous anhydride, S0 2 ; phosphoric anhydride, P 2 5 ; 
carbonic anhydride, C0 2 . When dissolved in water these 
oxides are converted into acids as represented in these 
equations : 

SO s + H 2 = H 2 S0 4 ; 
S0 2 + H 2 = H 2 SO s ; 
PA + H 2 = 3HP0 3 ; 
00 2 + H 2 = H 2 C0 3 . 

Silicic anhydride, Si0 2 , is an example of an acidic oxide 
which does not dissolve in water, but which does form 
salts when treated with basic hydroxides : 

Si0 2 + 2KOH = K 2 Si0 3 + H 2 0. 

As examples of basic oxides or oxides which when treated 
with water yield bases, the following may be taken : Calcium 
oxide, CaO; potassium oxide, K 2 0; barium oxide, BaO. 



PERIODIC ACID. 195 

As has already been shown, when treated with water 
these are respectively converted into calcium hydroxide, 
Ca(OH),; potassium hydroxide, KOH; and barium hy- 
droxide, Ba(OH) 2 . There are, however, many basic oxides 
which do not dissolve in water, but which, nevertheless, 
have the power to neutralize acids and form salts. This 
is true, for example, of aluminium oxide, A 2 10 3 ; lead 
oxide, PbO; manganous oxide, MnO; cupric oxide, CuO, 
etc. The action of such oxides upon acids takes place as 
represented below: 

ALA + 3H,S0 4 - A1 2 (S0 4 ) 3 + 3H 2 G; 
PbO + 2HN0 3 = Pb(NCV 2 -f H 2 0; 
MnO + 2HC1 ° = MnCl 2 + H 2 0; 
CuO + H 2 SO, = OuS0 4 + H 2 0. 

Periodic Acid, H 5 I0 6 . — This acid is analogous to per- 
chloric acid. Its salts are formed by oxidation of iodates 
or by heating iodates, just as perchlorates are formed by 
heating chlorates. The simplest way to make a periodate 
is to pass chlorine into a solution containing sodium 
hydroxide and sodium iodate, when a reaction takes place 
which is at least partly represented by the following equa- 
tion: 

NaI0 3 + 3NaOH + 01, = Na 2 H 3 IQ 6 + 2NaCl. 

The salt Na 2 H 3 I0 6 is difficultly soluble in water, and 
therefore separates from the solution. From the sodium 
salt the corresponding silver salt, Ag 2 H 3 I0 6 , can be 
obtained, and when this silver salt is treated with nitric 
acid it is converted into the simpler salt, AgI0 4 , which is 
evidently derived from the simpler acid, HI0 4 : 

2Ag 2 H 3 I0 6 + 2HN0 3 = 2AgN0 3 + 4H 2 + 2AgI0 4 . 

The acid when separated from its solutions is a crystallized 
solid which has the composition H 5 T0 G . When heated it 



196 COLLEGE CHEMISTRY. 

undergoes decomposition, losing water and oxygen, and 
yielding iodic acid. It cannot, however, be converted into 
a compound of the composition HI0 4 , for the loss of water 
is always accompanied by a loss of oxygen. Like iodic 
acid, periodic acid is a good oxidizing agent in conse- 
quence of the ease with which it gives up its oxygen. 

Compounds of Iodine with Chlorine. — When chlorine is 
passed over dry crystallized iodine it is absorbed, and a 
compound of the formula 101 is formed. This is a thick 
reddish-brown, very volatile liquid. Under proper condi- 
tions it solidifies in crystals, of which two kinds with 
different melting-points, 14° and 27°, have been described. 
Iodine chloride is decomposed by water, the products being 
iodic acid, hydrochloric acid, and free iodine. 

If the passage of chlorine over iodine is continued 
beyond the point required for the formation of the simple 
compound 101, the trichloride IC1 3 is formed. This is a 
crystallized compound of a yellow color. When heated it 
breaks down into chlorine and iodine monochloride. 
When treated with water it is partly dissolved without 
decomposition, but it is partly decomposed, yielding iodic 
acid, iodine monochloride, and hydrochloric acid. 

Compound of Iodine with Bromine. — There is only one 
compound of iodine and bromine known, and that is the 
one having the formula IBr. It is a crystallized compound 
which is formed by direct combination of the two elements. 
It is decomposed by heat and by water. 



Fluorine, F (At. Wt. 19). 

Occurrence. — This element occurs in large quantity in 
nature, and is widely distributed, but it is always in com- 
bination with other elements. It is found chiefly in com- 
bination with calcium, as fluor-spar or calcium fluoride, 
CaF 2 , and in combination with sodium and aluminium, as 
cryolite, a mineral which occurs abundantly in Greenland 
and has the composition represented by the formula 



FLUORINE. 197 

Na 3 AlF 6 or AlF 3 .3NaF. It is called fluorine from the fact 
that it occurs in fluor-spar, which in turn receives its name 
for the reason that it melts when heated and is therefore 
used as a flux in heating chemical substances together 
(from fiuo, I flow). On account of the remarkable power 
of combination of fluorine with other elements, all attempts 
to prepare it in the free condition failed until a few years 
ago, when its isolation was effected by passing an electric 
current through a solution of acid potassium fluoride, 
KHF 2 , in dry hydrofluoric acid contained in a platinum 
vessel. It has since been shown that a copper vessel may 
be used. 

Properties. — Fluorine is a light greenish-yellow gas, 
that has recently been converted into a liquid at a very low 
temperature. It acts upon almost all substances. Thus, 
it decomposes water, yielding ozone and hydrofluoric acid; 
it combines directly with hydrogen at the ordinary tem- 
perature; and with sulphur, phosphorus, iron, etc., with 
evolution of light and heat. It does not, however, act 
upon platinum. Owing to its active properties it is of 
course a difficult matter to isolate and preserve it. 

Hydrofluoric Acid, HF. — Hydrofluoric acid is made by 
treating a fluoride with sulphuric acid. Thus, when 
calcium fluoride or fluor-spar is used, this reaction takes 
place : 

CaF 2 + H 2 S0 4 = CaS0 4 + 2HF. 

The reaction must be performed in vessels of platinum or 
lead, as glass is disintegrated by the acid. In perfectly 
pure anhydrous condition it can be obtained by heating the 
pure dry salt KHF 2 , known as acid potassium fluoride. 
It is a liquid which boils at 19.4° and does not solidify 
even at a very low temperature. The pure dry acid in the 
liquid form does not act upon glass. It does not dissolve 
the acid-forming elements, but does dissolve most of the 
base-forming elements with evolution of hydrogen and 
formation of fluorides. The gas acts upon the skin, 



198 COLLEGE CHEMISTRY. 

causing swellings and violent pains. Inhaled it is poison- 
ous. To preserve it, vessels of platinum or caoutchouc 
must be used. In the moist condition it attacks glass, 
converting the silicon into the fluoride, SiF 4 , and the 
metals into their fluorides. A silicate of the formula 
CaSiO s would undergo the changes represented in the fol- 
lowing equation : 

CaSi0 3 + 6HF = CaF 2 + SiF 4 + 3H 2 0. 

Silicon fluoride is a gas, and calcium fluoride is soluble in 
acids. Thus calcium silicate, which is insoluble in water, 
is so changed by hydrofluoric acid as to be rendered solu- 
ble. In a similar way glass, which is a compound resem- 
bling calcium silicate, is rendered soluble, or is, as we 
commonly say, dissolved, by hydrofluoric acid. 

When an aqueous solution of hydrofluoric acid is boiled 
it passes over at 120°, and the distillate contains 36 to 38 
per cent of the acid. 

Hydrofluoric acid is used for the purpose of etching 
glass, particularly for marking scales on thermometers and 
other graduated glass instruments. The glass is covered 
with a thin layer of wax or paraffin and, at the places 
where the etching is wanted, marks are made through the 
paraffin, so that the glass is exposed. Those parts of the 
glass which are covered are not acted upon by the hydro- 
fluoric acid, while those parts which are not covered are 
corroded and, when the paraffin is removed, permanent 
marks are found corresponding to those made through the 
paraffin. A solution of hydrofluoric acid in water is 
manufactured and sold in rubber bottles. 

The specific gravity of hydrofluoric acid gas at about 
100° leads to the molecular weight corresponding to the 
formula HF, fluorine having the atomic weight 19. At 
about 30° the specific gravity corresponds to the formula 
H 2 F 2 . At lower temperatures the molecular weight 
appears to be still greater. 



MANGANESE. 199 

Family VII, Group A — Manganese. 

Manganese belongs to the same family as the halogens, 
and resembles them in some respects, but at the same time 
it differs from them quite markedly in other respects. It 
acts in fact in two different ways, and is one of those 
elements, already referred to, that are both acid-forming 
and base-forming. Some of its compounds with hydrogen 
and oxygen are distinctly acid, others are distinctly basic. 
So far as it acts like the members of the chlorine family a 
brief reference to it here is desirable. On the other hand, 
it will be dealt with chiefly in connection with those base- 
forming elements which it most resembles, as, for ex- 
ample, iron. 

Manganese occurs in nature principally in the form of 
pyrolusite or manganese dioxide, Mn0 2 , also known as the 
black oxide of manganese. It forms with oxygen com- 
pounds of the following formulas: MnO, Mn 2 3 , Mn 3 4 , 
Mn0 2 , and Mn 2 7 . When a compound of manganese is 
subjected to the influence of powerful oxidizing agents in 
the presence of an alkali it is converted into a salt of 
manganic acid, H 2 Mn0 4 , which in its composition resem- 
bles sulphuric acid. If the salt of manganic acid thus 
obtained is dissolved in water it undergoes partial decom- 
position, which is complete if the solution is boiled, or if 
carbon dioxide is passed through it. The change consists 
in the transformation of manganic acid into permanganic 
acid, HMn0 4 : 

3H 2 Mn0 4 = 2HMn0 4 + MnO., + 2H 2 0; or 
3K>ln0 4 + 2H 2 = 2KMn0 4 + MnO, + 4KOH. 

Permanganic acid, HMn0 4 , is a compound which in 
many respects resembles perchloric acid. It can be 
obtained in water solution by decomposition of certain of 
its salts, but like perchloric acid it is easily decomposed. 
In consequence of the ease with which it gives up oxygen 
it is a good oxidizing agent, and is extensively used in the 



200 COLLEGE CHEMISTRY. 

laboratory in this capacity. It is employed in the form 
of the potassium salt, potassium permanganate, KMn0 4 , 
which will receive special attention under the head of 
Manganese Compounds. In order, however, to make 
clear the difference in conduct between perchloric and 
permanganic acids a few characteristic facts will be men- 
tioned here. The conduct of permanganic acid and of 
potassium permanganate will be understood, if it is borne 
in mind that in the presence of substances of strongly acid 
character manganese tends to act as a base-forming ele- 
ment, and in this capacity to form salts with the acids. 
Thus in presence of sulphuric acid potassium permanganate 
forms potassium sulphate, manganous sulphate, and 
oxygen, if there is anything present which has the power 
to take up oxygen. In the salts in which it plays the part 
of a metal manganese is generally bivalent. With hydro- 
chloric acid, as we have already seen in studying the action 
of hydrochloric acid upon manganese dioxide, it forms the 
chloride MnCl 2 . When now potassium permanganate is 
treated with hydrochloric acid it is decomposed according 
to the following equation : 

2KMn0 4 + 6H01 = 2KC1 + 2MnCl 2 + 3H 2 + 50. 

Similarly, with sulphuric acid manganous sulphate is 
formed, thus: 

2KMn0 4 + 3H 2 S0 4 = K 2 S0 4 + 2MnS0 4 + 3H 2 + 50. 

Such reactions do not take place with perchloric acid, as 
chlorine is entirely lacking in the power to enter into acids 
in the place of the hydrogen and form salts. 

Manganese forms some other acids besides permanganic 
acid, but they exhibit little or no- analogy with compounds 
of chlorine, and their study will therefore be postponed 
until manganese is taken up. The point of chief interest 
to be noted here is that this element is unmistakably like 
chlorine in its highest oxygen compounds, but entirely 



EXPERIMENTS WITH BRQMJN-E„ETC 201 

; ; * 

different from it in most of its compounds. Tlie com- 
pound manganese heptoxide, Mn 2 7? stands in the relation 
of an anhydride to permanganic acid. In water solution 
it passes over into the acid: 

Mn 2 7 + II 2 = 2HMn0 4 . 

~t is formed by treating potassium permanganate with 
.iie most concentrated sulphuric acid: 

2KMn0 4 + H 2 S0 4 = K 2 S0 4 + Mn 2 7 + H 2 0. 

It is extremely unstable, giving up oxygen readily. In 
contact with organic substances or other substances which 
have the power to take up oxygen it decomposes so rapidly 
as frequently to lead to explosions. It is plainly analogous 
to the oxide of chlorine of the formula C1 2 7 . 

EXPERIMENTS. 

Preparation" of Bromine. 

Experiment 81. — Mix together 3.5 grams potassium bromide 
and 7 grams manganese dioxide. Put the mixture into a 500 cc. 
flask ; connect with a condenser (see Fig. 28). Mix 15 cc. con- 
centrated sulphuric acid and 90 cc. water. After cooling pour 
the liquid on the mixture in the flask. Gently heat, when bromine 
will be given off in the form of vapor. A part of this will con- 
dense and collect in the receiver. Perform this experiment under 
a hood with a good draught. 

Hydrobromic Acid. 

Experiment 82. — In a small porcelain evaporating-dish put a 
few crystals of potassium bromide. Pour on them a few drops of 
concentrated sulphuric acid. The white fumes of hydrobromic 
acid and the reddish-brown vapor of bromine are noticed. Treat 
a few crystals of potassium or sodium chloride in the same way. 
What difference is there between the two cases ? 

The preparation of hydrobromic acid may be shown in the lec- 
ture-room as follows : 



202 



COLLEGE CHEMISTRY. 



Experiment 83. — Arrange an apparatus as shown in Fig. 40. 
In the flask put 1 part red phosphorus and 2 parts water. Let 




Fig. 40. 

10 parts bromine gradually drop into the flask from the glass- 
stoppered funnel. Pass the gas through a U-tube loosely packed 
with asbestos containing moistened red phosphorus in order to 
free the hydrobromic acid from bromine, which to some extent 
passes over with it. Collect some of the gas in water, and examine 
the solution. How does the gas act when allowed to escape in 
the air ? Fill a cylinder with the gas in the same way as was done 
with hydrochloric acid, and fill another with chlorine. While 
covered with glass plates bring their mouths together. Then with- 
draw the plates. "What change is observed ? What is this due to ? 
Experiment 84.— To a dilute solution of sodium hydroxide add 
bromine water made by shaking up a little liquid bromine in a 
bottle with water. What change takes place ? Add sulphuric acid 
until the liquid shows an acid reaction. What takes place ? The 
changes here referred to are perfectly analogous to those which 
would take place if chlorine were used instead of bromine. Shake 
a solution containing a little free bromine with ether ; with chlor- 
oform ; with carbon disulphide. What changes do you observe ? 

Iodide. 



Experiment 85. — Mix about 2 grams of sodium or potassium 
iodide and 4 grams manganese dioxide. Treat with a little con- 
centrated sulphuric acid in a one- to two-litre flask. Heat gently 



EXPERIMENTS IVITH IODINE, ETC. 203 

on a sand-bath. Gradually the vessel will be filled with the beau- 
tiful colored vapor of iodine. In the upper parts of the flask 
some of the iodine will be deposited in the form of crystals of a 
grayish-black color. 

Experiment 86. — Make solutions of iodine in water, in alcohol, 
and in a water solution of potassium iodide. Use small quantities 
in test-tubes. 

Experiment 87. — Dissolve a piece of potassium iodide the size 
of a small pea in about 100 cc. water in a stoppered cylinder. 
Add enough carbon disulphide to make a layer about an inch 
thick at the bottom of the cylinder. Shake the two liquids to- 
gether. Does the carbon disulphide become colored ? Add a 
drop of chlorine water and shake again. What difference do you 
observe in the two cases? Explain this. Try the same experi- 
ment, using chloroform instead of carbon disulphide. 



Iodine can be detected by Means of its Action upon 
Starch-paste. 

Experiment 88. — Make some starch-paste by covering a few 
grains of starch in a porcelain evaporating dish with cold water, 
grinding this to a paste, and pouring 200-300 cc. boiling-hot 
water on it. After cooling add a little of this paste to a dilute 
water solution of iodine in a solution of potassium iodide in water. 
The solution will turn blue. Now add a little of the paste to a 
dilute water solution of potassium iodide alone. Is there any 
change ? Add a drop or two of a solution of chlorine in water. 
Why the difference ? Will not chlorine water alone act this way 
toward starch-paste ? 

Action of Sulphuric Acid upon Potassium Iodide. 

Experiment 89. — Bring a piece of potassium iodide the size of 
a pea in a dry test-tube ; add one drop of water and three or four 
drops of concentrated sulphuric acid ; the salt becomes brown ; 
heat gently ; violet-colored vapor escapes, and with it a gas with an 
odor like that of rotten eggs. At the same time a yellow coating 
appears on the inside of the tube above the acid. Add five or six 
drops more of the acid and continue to heat gently. The bad 
odor first noticed disappears gradually, and another, quite differ- 
ent odor, irritating to the throat, is now perceptible. This is sul- 
phur dioxide, SO a . 



204 COLLEGE CHEMISTRY. 



Iodic Acid. 

Experiment 90.— Pass chlorine into a test-tube containing 
iodine in suspension in water ; or add chlorine water. What be- 
comes of the iodine ? 

Experiment 91. — Add chlorine water to a dilute solution of 
potassium iodide, and note the successive changes. 

Experiment 92.— Dissolve iodine in caustic soda. Add an acid 
to the solution. Explain the changes. 

Hydrofluoric Acid. 

Experiment 93.— In a lead or platinum vessel put a few grams 
(5-6) of powdered fluor-spar and pour on it enough concentrated 
sulphuric acid to make a thick paste. Cover the surface of a 
piece of glass with a thin layer of wax or paraffin, and through 
this scratch some letters or figures, so as to leave the glass ex- 
posed where the scratches are made. Put the glass over the ves- 
sel containing the fluor-spar, and let it stand for some hours. 
Take off the glass, scrape off the coating, and the figures which 
were marked through the wax or paraffin will be found etched on 
the glass. 



CHAPTEE XIII. 

THE ELEMENTS OF FAMILY VI, GROUP B: 
SULPHUR.— SELENIUM.— TELLURIUM. 

Introductory. — The elements of this group bear to 
oxygen a relation somewhat similar to that which the ele- 
ments of Group B, Family VII, bear to fluorine. The 
three members sulphur, selenium, and tellurium resemble 
one another fully as strikingly as chlorine, bromine, and 
iodine do. Their compounds bear a general resemblance 
to those of oxygen, and yet they form very characteristic 
compounds with oxygen, while oxygen forms no analogous 
compounds with any of them. Just as iodine forms a 
compound with fluorine of the formula IF. , but fluorine 
does not form with iodine a compound FI 5 , so sulphur, 
selenium, and tellurium form with oxygen the compounds 
S0 3 , Se0 3 , and Te0 3 , while oxygen does not form 
analogous compounds with these other elements. The 
valence of the elements of this group towards hydrogen is 
2, as shown in the compounds H 2 0, H 2 S, H 2 Se, and H 2 Te. 
Of oxygen and sulphur there are other hydrogen com- 
pounds, as hydrogen dioxide, H 2 2 , and an analogous 
compound of sulphur, but there is much uncertainty in 
regard to the constitution of the latter compound, and 
practically nothing is known in regard to the valence of 
sulphur in it. In general the hydrogen valence is con- 
stant. Towards the members of the chlorine group the 
valence varies from 2 to 6 Oxygen never exhibits a 
higher valence than 2 towards chlorine and its analogues. 
The compound OCl 2 illustrates the bivalence of oxygen 
towards chlorine. Sulphur forms with chlorine the com- 
pounds S 3 C1 2 and SC1 2 , which are analogous to the hy- 

205 



206 COLLEGE CHEMISTRY. 

drogeu compounds H 2 S 2 and H 2 S, and in both of them the 
sulphur is probably bivalent. It also forms the compound 
SC1 4 , in which it is quadrivalent. With iodine it is said 
to form the unstable compounds S 2 I 2 and SI 6 . With 
fluorine it forms the stable fluoride, SF 6 . Selenium and 
tellurium form similar compounds, and, in general, the 
stability of the compounds of the members of the sulphur 
group with the members of the chlorine group increases 
in the order sulphur, selenium, tellurium. 

Towards oxygen the three elements of the sulphur 
group are quadrivalent and sexivalent, as seen in the com- 
pounds S0 2 , Se0 2 , Te0 2 , and S0 3 , Se0 3 , and Te0 3 . 

Of the three elements of this group sulphur occurs in 
greatest abundance in nature, selenium next in order, and 
finally tellurium. Just as bromine frequently accompanies 
chlorine, so selenium frequently accompanies sulphur, but 
it is always present in much smaller quantity than sulphur. 
Tellurium occurs in very small quantity relatively, and not 
uncommonly in combination with valuable metals like gold 
and silver. Large quantities of sulphur are found in the 
native or uncombined condition. Only extremely small 
quantities of selenium and tellurium are found native. 

Sulphur, S (At. Wt. 32.06). 

Occurrence. — The principal deposits of native sulphur 
are found in Sicily, Italy, and Spain. In California also 
there is a considerable deposit. In general, sulphur is 
found near dying or extinct volcanoes. Sulphur occurs in 
nature, further, in the form of the hydrogen compound, 
hydrogen sulphide, H 2 S, issuing from the earth in volcanic 
regions, and in solution in some natural waters, known as 
" sulphur waters." The oxide S0 2 is likewise found 
issuing from the earth in volcanic regions. Compounds 
of sulphur with metallic elements, as with iron, copper, 
lead, zinc, are very abundant. Such compounds are iron 
pyrites, FeS 2 ; copper pyrites, CuFeS 2 ; galenite, PbS; and 
zinc blende, ZnS. Some sulphates are widely distributed 



SULPHUR. 207 

and occur in large quantities; for example, gypsum or 
calcium sulphate, CaS0 4 + 2H 2 0; barium sulphate, or 
heavy spar, BaSOj lead sulphate, PbS0 4 . Finally, sul- 
phur occurs in a few animal and vegetable products in 
combination with carbon, hydrogen, and, generally, with 
nitrogen. 

Extraction of Sulphur from its Ores. — By far the largest 
Quantity of sulphur found in the market is taken from the 
mines in Sicily. Of these mines there are between 250 
and 300. When taken from the mines it is mixed with 
many earthy substances, from which it must be separated. 
In Sicily the separation is generally effected by piling the 
ore so as to leave passages for air, covering the piles with 
earthy matter to prevent free access of air and then setting 
fire to them. A part of the sulphur burns, and this causes 
the rest of it to melt. The molten sulphur runs down to 
the bottom of the pile, and by a proper arrangement it is 
drawn off from time to time. If the pile of ore were not 
covered up, and the air were allowed to enter freely, the 
sulphur would burn up, and be converted into the gas 
sulphur dioxide. The " crude brimstone " obtained in the 
manner described is afterwards refined by distillation, and 
it is this refined or distilled sulphur which is met with 
in the market under the names "roll brimstone" and 
"flowers of sulphur." 

The distillation is carried on in retorts made of earthen- 
ware, and these are connected with large chambers of 
brick- work. When the vapor of sulphur first comes into 
the conden sing-chamber it is suddenly cooled, and hence 
deposited in the form of a fine powder. This is called 
" flowers of sulphur." After the distillation has continued 
for some time the vapor condenses in the form of a liquid, 
which collects at the bottom of the chamber. This is 
drawn off into slightly conical wooden moulds, and takes 
the form of "roll brimstone" or "stick sulphur." 

Some sulphur is obtained from iron pyrites by heating 
in closed vessels. The change which takes place on heat- 
ing iron pyrites is perfectly analogous to ttat which takes 



2o8 COLLEGE CHEMISTRY. 

place on heating manganese dioxide in preparing oxygen. 
A sulphur compound of the formula Fe 3 S 4 and free sulphur 
are formed in the former case, as the compound of man- 
ganese and oxygen, MngO^ , and free oxygen are formed in 
the latter: 

3FeS 2 = Fe 3 S 4 + S 2 ; 
3Mn0 2 = Mn 3 4 + 2 . 

Properties. — Sulphur is a yellow, brittle substance which 
at — 50° is almost colorless. If heated rapidly it melts 
at 114.5°, forming a thin, straw-colored liquid. When 
heated to a higher temperature it becomes darker in color, 
and at 200° to 250° it is so viscid that the vessel in which 
it is contained may be turned upside down without danger 
of its running out. Finally, at 448.4° it boils, and is then 
converted into a brownish-yellow vapor. When molten 
sulphur solidifies, or when it is deposited from a solution, 
it takes the form of crystals. But, strange to say, the 
crystals formed from molten sulphur are entirely different 
from those deposited from cold solutions of sulphur. The 
latter belong to the rhombic system. They are octahedrons 
with a rhombic base. And this is the form of the sulphur 
found in nature. The former are honey-yellow needles. 
An examination of these needles shows that the angles 
which their faces form with one another are not the same 
as the angles formed by the faces of the octahedrons, and 
that they belong to an entirely different system — the 
monoclinic. 

The rhombic crystals of sulphur can be made by dissolv- 
ing "roll brimstone" in carbon disulphide and allowing 
the solution to stand. When the liquid has sufficiently 
evaporated, the sulphur will appear in larger or smaller 
rhombic crystals, according to the conditions. A com- 
parison of the crystals thus obtained with natural crystals 
will show that the two have identical or very similar forms. 
The formation of the needles or monoclinic crystals may 
be shown by melting a considerable quantity, say a pound 



SULPHUR. 209 

or two, of roll brimstone in a sand or Hessian crucible, 
and allowing the liquid mass to cool slowly. When a thin 
crust has formed on the surface this should be perforated, 
and the remainder of the liquid sulphur poured off. The 

crucible will then be found lined with long, dark-yellow, 
lustrous needles which do not look at all like those obtained 
from the solution in carbon disulphide. If the monoclinic 
needles are allowed to lie unmolested they gradually 
undergo change spontaneously. They lose their lustre and 
become lighter in color; and now they consist o\' minute 
crystals like those found in nature. They have changed 
to the rhombic form. It is evident, therefore, that the 
arrangement of the particles in the monoclinic crystals of 
sulphur is not a stable one at ordinary temperatures, 
whereas it has been shown that it is stable at temperatures 
above 100°. The change from the monoclinic to the 
rhombic form is accompanied by a considerable evolution 
of heat. Monoclinic sulphur melts at 120°. 

Substances which crystallize in two distinct forms are 
called dimorphous. We shall see that carbon also crystal- 
lizes in two different forms, and that this kind of relation 
is not unfrequently met with among chemical compounds. 
The difference between the two varieties of sulphur sug- 
gests that observed between the two forms of oxygen. 
Whether the explanation is the same in the two cases is 
doubtful. 

Sulphur can also be obtained in the amorphous, or un- 
crystallized, condition. If sulphur is heated considerably 
above its melting-point and then quickly cooled under 
water it remains for some time soft and dough-like, and 
while in this condition it is amorphous. If allowed to 
stand it gradually becomes hard and brittle. 

When separated from certain compounds which are in 
solution in water, the sulphur is in a very finely-divided 
condition, and gives the liquid the appearance of milk. 
This i"s seen on adding hydrochloric acid to a solution of 
sodium thiosulphate, or hyposulphite, as it is generally 
called. 



210 COLLEGE CHEMISTRY. 

On treating certain varieties of sulphur with carbon 
disulphide they are found to dissolve completely. This is 
true, for example, of the natural crystals and of those 
made artificially by deposition from a solution in carbon 
disulphide. On the other hand, sulphur in the form of 
"flowers of sulphur" is only partly soluble in the liquid. 
There are therefore two forms of sulphur to be distin- 
guished between, the soluble and the insoluble. The cause 
of the difference between these modifications is not known. 
"Stick sulphur" is mostly so]uble, while in the "flowers 
of sulphur " there is at times a considerable percentage of 
the insoluble variety. Sulphur is insoluble in water, and 
slightly soluble in alcohol and ether. 

Sulphur is a much less active element chemically than 
the members of the chlorine group, and also less active 
than oxygen. Generally speaking, however, it conducts 
itself like oxygen. It combines directly though not easily 
with hydrogen, and it combines readily with most metals, 
forming compounds called sulphides which, so far as their 
composition is concerned, are analogous to the oxides. 
Thus when heated together with iron, copper, or lead, 
combination takes place readily with evolution of heat and 
light. 

In its chemical conduct it differs markedly from the 
members of the chlorine group. When heated to a suffi- 
ciently high temperature in the air or in oxygen, sulphur 
forms the compound sulphur dioxide, S0 2 , and, under 
certain conditions which will be described farther on, this 
combines with more oxygen to form the trioxide S0 3 . 
Further, it combines with nearly all the acid-forming ele- 
ments if heated with them to a sufficiently high tempera- 
ture. It combines with most other elements less readily 
than oxygen and it forms less stable compounds. Thus, 
its compound with hydrogen is decomposed much more 
readily into its elements than water is; and the sulphides 
or its compounds with metals are decomposed when they 
are heated in oxygen, the oxygen displacing the sulphur, 
much as chlorine displaces bromine; though there is a 



SULPHUR. 211 

difference between the two cases to be found in the fact 
that sulphur itself unites readily with oxygen, and this 
facilitates the decomposition of the sulphides by the action 
of oxygen. 

When treated with powerful oxidizing agents in the 
presence of water sulphur is converted into sulphuric acid. 

The specific gravity of the vapor of sulphur varies with 
the temperature, and is such as to indicate that at tem- 
peratures not far above the boiling-point the molecule con- 
sists of eight atoms, while at the temperature 800° and 
higher the molecule consists of two atoms. 

Uses of Sulphur, — Enormous quantities of sulphur are 
used in the manufacture of sulphuric acid, and of gun- 
powder. It is also used in the manufacture of fireworks 
of various kinds. Burning sulphur gives sulphur dioxide, 
which is extensively employed for bleaching wool, silk, 
straw, etc. "When caoutchouc is thoroughly mixed with 
sulphur or some sulphur compound it becomes vulcanized. 



Compounds of Sulphur with Hydrogen. 

The principal compound of sulphur and hydrogen is 
analogous in composition to water. It is known as 
hydrogen sulphide or sulphuretted hydrogen, and has the 
formula H 2 S. Besides this there is at least one other 
compound which contains a larger proportion of sulphur, 
and probably has the composition H 2 S 2 . Farther, there 
are reasons for supposing that still more complex com- 
pounds can exist, but owing to their instability it has been 
found impossible to isolate them. 

Hydrogen Sulphide, Sulphuretted Hydrogen, H 2 S. — 
When hydrogen is passed over highly-heated sulphur the 
two elements combine to form hydrogen sulphide. The 
action is, however, quite incomplete and is not to be com- 
pared with that which takes place when hydrogen and 
oxygen are heated together. This compound of sulphur 
and hydrogen occurs in nature in solution in the so-called 



212 COLLEGE CHEMISTRY. 

" sulphur waters," which are met with in many parts of 
the earth. It is formed by heating organic substances 
which contain sulphur, just as water is formed by heating 
organic substances which contain oxygen. It is formed, 
further, by decomposition of organic substances which 
contain sulphur, as, for example, the albumen of eggs. 
The odor of rotten eggs is partly due to the formation of 
hydrogen sulphide. It is formed by the action of acids 
upon sulphides or hydrosulphides, just as water is formed 
by the action of acids upon oxides or hydroxides. Thus 
hydrochloric acid and ferrous sulphide, FeS, give ferrous 
chloride, FeCl 2 , and hydrogen sulphide: 

FeS + 2H01 = Fe01 2 + H 2 S; 

just as ferrous oxide, FeO, and hydrochloric acid give 
ferrous chloride and water : 

FeO + 2HC1 = FeCl 2 + H 2 0. 

In the laboratory, where the gas is extensively used, it 
is generally prepared from ferrous sulphide, FeS, and 
dilute sulphuric acid, which are brought together at the 
ordinary temperature in a flask such as is used in making 
hydrogen. The reaction is like that which takes place 
between ferrous sulphide and hydrochloric acid. It is 
represented by this equation : 

FeS + H 2 S0 4 =± FeS0 4 + H 2 S. 

Properties, — Hydrogen sulphide is a colorless, trans- 
parent gas of the specific gravity 1.178. It has an ex- 
tremely disagreeable odor, somewhat suggestive of that of 
rotten eggs. It is poisonous, even small quantities causing 
headache, vertigo, nausea, and other bad symptoms. It 
burns with a blue flame, forming water and sulphur 
dioxide : 

H 2 S + 30 = H 2 + S0 2 . 

If, however, the air has not free access, as when the gas is 



COMPOUNDS OF SULPHUR WITH HYDROGEN. 213 

burned in a cylinder open at one end, only a part of the 
sulphur is burned, the rest being deposited upon the walls 
of the vessel, while the hydrogen burns. The gas is solu- 
ble in water, about three volumes being taken up at 
ordinary temperatures. This solution is used to some 
extent in the laboratory instead of the gas, but, owing to 
the fact that it undergoes change in consequence of the 
action of the oxygen of the air, it is not as valuable as it 
would be if it were more stable. The change consists 
simply in the oxidation of the hydrogen and the separation 
of the sulphur. If a bottle containing a solution of 
hydrogen sulphide is allowed to stand for a few days, par- 
ticularly if it is opened from time to time, the odor of the 
gas will disappear and a deposit of sulphur will be noticed 
in the bottle. The liquid is then nothing but water. 
When the solution is boiled it loses all the gas contained 
in it. 

Hydrogen sulphide is easily decomposed into its ele- 
ments. It requires a temperature of only a little above 
400° to effect direct decomposition. In consequence of 
this instability it causes a number of changes which the 
analogous compound water cannot effect. The relations 
here are similar to those which exist between hydrochloric 
and hydriodic acids. Hydrochloric acid is very stable, 
while hydriodic acid breaks down readily into hydrogen 
and iodine. Therefore hydriodic acid, as we have seen, 
acts as a reducing agent, while hydrochloric acid does not. 
So, also, hydrogen sulphide acts as a reducing agent. 
Thus, if it is passed into concentrated sulphuric acid this 
reaction takes place : 

H 2 S0 4 + H 2 S = 2H 2 + S + S0 2 . 

The action is to be traced to the decomposition of the 
hydrogen sulphide into hydrogen and free sulphur, the 
hydrogen then acting upon the sulphuric acid thus: 

H 2 S0 4 + 2H = 2H 2 + S0 2 . 



214 COLLEGE CHEMISTRY. 

With hydriodic acid the reduction may go farther, as has 
been seen; with hydrobromic acid, however, the action 
takes place practically in the same way as with hydrogen 
sulphide. 

Chlorine, bromine, and iodine act upon hydrogen sul- 
phide, setting the sulphur free and uniting with the 
hydrogen : 

H 2 S + Cl 2 = 2HC1 + S. 

The instability of hydrogen sulphide is further shown 
by the ease with which it is decomposed by metals with 
liberation of hydrogen and formation of sulphides. It 
will be remembered that at high temperatures several 
metals decompose water, but that at ordinary temperatures 
only a few decompose it easily. Hydrogen sulphide acts 
much more readily; a number of metals which do not act 
upon water even at high temperatures, as silver, gold, and 
mercury, decompose this gas at ordinary temperatures. 

Hydrogen sulphide acts upon metallic oxides, converting 
them into sulphides, as for example : 

CuO + H 2 S = CuS + H 2 0. 

Action of Hydrogen Sulphide upon Solutions of Salts — 
Use in Chemical Analysis. — Hydrogen sulphide is exten- 
sively used in every chemical laboratory as a reagent in 
chemical analysis. In order that its action may be under- 
stood a few words of explanation are necessary. Sulphur, 
as we have seen, readily unites with the metallic or base- 
forming elements, forming with them the sulphides. Some 
of these are insoluble in water and in dilute acids; others 
are soluble in dilute acids but insoluble in water and in 
ammonia; while still others are soluble even in water. 
This plainly furnishes a basis for the classification of the 
base-forming elements into three groups for the purpose 
of analysis. We have : 

I. Metals whose sulphides are insoluble in water and are 
not decomposed by dilute acids. This is called the hy- 






HYDROGEN SULPHIDE. 215 

drogen sulphide group. It includes lead, bismuth, silver, 
mercury, copper, cadmium, gold, platinum, tin, antimony, 
and arsenic. 

II. Metals whose sulphides are insoluble in water but 
are decomposed by dilute acids. They are therefore not 
precipitated by hydrogen sulphide, but are precipitated by 
soluble sulphides. As ammonium sulphide is used for the 
purpose of precipitating these, the group is known as the 
ammonium sulphide group. It includes iron, nickel, 
cobalt, manganese, thallium, zinc, and uranium. Further, 
the two elements aluminium and chromium are thrown 
down with the above, not as sulphides but as hydroxides, 
and they are therefore included in the group. 

III. Metals whose sulphides are soluble in water. This 
group includes all the metals not included in the above two. 

By taking advantage, then, of the properties of the sul- 
phides of the metals they can be divided into these three 
groups, and the detection of any particular element is thus 
facilitated. If hydrogen sulphide is passed through a 
solution, and a precipitate formed, we know that this can 
contain only those metals which belong to the hydrogen 
sulphide group; and so on. Now, if the precipitate 
formed with hydrogen sulphide is treated with certain 
other reagents other changes take place, and by further 
study it is quite possible, and indeed comparatively simple, 
to determine which of the members of the group are 
present. 

According to the theory of ions (see pp. 90, 161), hy- 
drogen sulphide is dissociated in aqueous solution, forming 
the ions HS and H, and also S and HH. So, also, the 
metallic salts are dissociated. For example, silver nitrate, 
AgNO, , in aqueous solution gives the ions Ag and N0 3 . 
Now when hydrogen sulphide is passed into a solution of 
silver nitrate the reaction that takes place is to be repre- 
sented thus : 



2(Ag + N0 3 ) + S + HH = Ag 2 S + 2H + 2N0 3 . 



2l6 COLLEGE CHEMISTRY. 

The silver sulphide is removed from the solution and i? 
obtained as a solid precipitate, while the ions of nitric acid 
are in the solution, forming, as we say, a dilute solution 
of nitric acid. This dilute nitric acid does not act upon 
the silver sulphide, so that the reaction is complete, all the 
silver ions being carried out of the solution by the sulphur 
ions provided enough hydrogen sulphide is passed into the 
solution. 

When, however, hydrogen sulphide is passed into a 
solution of ferrous chloride, FeCl 2 , the results are differ- 
ent. In this case, we should naturally expect to find thai 
the reaction would take place as represented in the equa- 
tion: 

IV+ 2C1 + 2H + S = FeS + 211 + 20L 

But the ferrous sulphide, FeS, is acted vcf/M by hydro- 
chloric acid in solution, giving hydrogen &ylphide and 
ferrous chloride, so that the change above represented 
cannot take place. 

If, however, a soluble sulphide is added to a solution of 
a salt of iron or of any metal that acts in the same way, 
the sulphide is precipitated. Thus, if instead of passing 
hydrogen sulphide, a solution of potassium sulphide is 
added, reaction takes place thus : 

p e + 201 + 2K + S = FeS + 2K + 201. 

In the resulting solution the hydrogen ions of hydrochloric 
acid are not present and the ions of potassium chloride 
which are present do not act upon the ferrous sulphide. 
Therefore the ferrous sulphide is precipitated. 

Hydrosulphides. — The action of hydrogen sulphide 
shows that it belongs to the class of acids. When it acts 
upon an oxide the corresponding sulphide and water are 
formed. But just as there are sulphides that are derived 
from hydrogen sulphide by the substitution of metallic 
elements for both hydrogen atoms, so there are hydrosul- 

r 



HYDROGEN PERSULPH1DE. 21 7 

phides that are derived from it by the replacement of only 
one of the two atoms of hydrogen in the molecule. The 
sulphides correspond to the oxides, and the hydrosulphides 
to the hydroxides. Thus the analogous oxygen and sul- 
phur compounds of potassium are : 

K 2 K 2 S 
KOH KSH. 



We speak of sulphides and hydrosulphides as salts of 
hydrogen sulphide. In consequence of this power to form 
salts in the same way in general that acids do, hydrogen 
sulphide is sometimes called sulpJiydric acid, and the salts 
of the formula MSH, sulphydrates. The name sulphydrate 
is analogous to hydrate, which, as has been pointed out, 
is used by some to designate the compounds of the formula 
MOH or the hydroxides. Between the acid, hydrogen 
sulphide, and the neutral compound, water, there is no 
fundamental difference. The difference is simply one of 
degree. In the present system of chemistry, which is 
largely an oxygen system, water is regarded as the con- 
necting link between acids and bases, as was shown on 
p. 160. But we might with equal right base our defini- 
tions and conceptions of acids and bases upon the conduct 
of sulphur compounds, and thus build up a sulphur 
sj^stem. In such a system hydrogen sulphide would be 
the connecting link between acids and bases. 

Hydrogen Persulphide, H 2 S 2 (?). — The sulphides of 
certain metals, particularly the so-called alkali metals, 
sodium and potassium, combine with sulphur to form the 
polysulphides, examples of which are K 2 S 2 , K 2 S 3 , K 2 S 4 , 
and K 2 S.. When these are decomposed with dilute acids, 
compounds of hydrogen and sulphur are formed. It has 
thus far, however, been impossible to determine whether 
more than one such compound is formed, for the reason 
that there is no means of deciding whether the substances 
formed are chemical compounds or mere mixtures of sul- 
phur and some one compound of sulphur and hydrogen. 



2i8 COLLEGE CHEMISTRY. 

Hydrogen persulphide is a liquid with a very disagreeable 
odor. Just as hydrogen dioxide decomposes readily into 
water and oxygen, so hydrogen persulphide decomposes 
readily into hydrogen sulphide and sulphur. 

Compounds of Sulphur with Members of the Chlorine 
Group. — Sulphur combines directly with chlorine and 
forms the compounds S 2 C1 2 , SC1 2 , and SC1 4 . Of these the 
first is the most stable. This can be boiled without under- 
going decomposition. The second, sulphur dichloride, 
SC1 2 , undergoes decomposition into chlorine and sulphur 
monochloride at the boiling-point; while sulphur tetra- 
chloride exists only at low temperatures. All these 
compounds are decomposed by water, yielding oxygen 
compounds. With iodine sulphur forms an unstable 
hexiodide, SI 6 : with fluorine, a very stable hexafluoride, 
SF 6 . 

Selenium, Se (At. Wt. 79.1). 

Occurrence. — Selenium occurs only in small quantity in 
nature. It was first found in the deposit formed in a sul- 
phuric-acid chamber (see p. 227), and owed its origin to 
the presence of small quantities of selenides in the sul- 
phides used in the operation. It was found to resemble 
tellurium, and for that reason was called selenium, from 
creXrjv)], the moon, tellurium receiving its name from 
tellus, the earth. Selenium occurs in a number of the 
natural sulphides, such as iron pyrites, copper pyrites, 
zinc blende, etc. ; and when these are treated in a current 
of air to decompose them and convert the sulphur into 
sulphur dioxide, the selenium is also oxidized, and the 
selenium dioxide thus formed is carried into the flues and 
other parts of the apparatus. As it is a solid it is easily 
condensed, and gradually a considerable quantity collects 
in the flues. This flue-dust is the best material from 
which to make selenium. 

Properties. — There are two modifications of selenium, 
corresponding to those of sulphur. One is soluble in 



TELLURIUM. 219 

carbon disulphide, the other is riot. The soluble form is 
obtained by reducing selenious acid by means of sulphurous 
acid, or by means of other reducing agents. The insoluble 
variety is obtained by melting selenium, then cooling it 
down suddenly to 210°, and keeping it at this temperature 
for some time. When it solidifies it is found to be no 
longer soluble in carbon disulphide. Selenium burns as 
sulphur does, forming the oxide Se0 2 , which has the odor 
of rotten horse-radishes. 



Tellurium, Te (At. Wt. 127). 

Occurrence. — Tellurium occurs in some gold ores in the 
native or uncombined condition, and also in combination 
with gold, silver, antimony, lead, and other metals. Tel- 
lurides occur, among other places in the United States, in 
California and Virginia. The general method of prepar- 
ing tellurium from its ores is the same as that by which 
selenium is made. The tellurium is oxidized to tellurious 
acid, H 2 Te0 3 , which is isolated, and then reduced by 
means of sulphurous acid. 

Properties. — Tellurium is silver-white, and crystallizes 
easily. Heated in the air it burns, forming a thick white 
cloud of tellurium dioxide, Te0 2 . Treated with sulphuric 
acid it is oxidized to tellurious acid, H 2 Te0 3 , the sul- 
phuric acid being reduced. Mtric acid oxidizes it likewise 
to tellurious acid. Melted together with potassium nitrate 
it is converted into potassium tellurate, K 2 Te0 4 . 



EXPERIMENTS. 

Properties of Sulphur. 

Experiment 94.— Distil about 10 grams roll sulphur from an 
ordinary glass retort. What changes in color and in condition 
take place ? Collect the liquid sulphur formed by the condensa- 
tion of the vapor in a beaker-glass containing cold water. 



2 20 COLLEGE CHEMISTRY. 

Experiment 95. — Treat some powdered roll sulphur with car- 
bon disulphide and filter. Does it all dissolve ? Try the same ex- 
periment with dowers of sulphur. Does this all dissolve ? Put 
the solutions together and allow to evaporate. Examine the crys- 
tals deposited. Compare them with some natural crystals of sul- 
phur. See whether one of the crystals will completely dissolve 
in carbon disulphide. 

Experiment 96. — In a covered sand or Hessian crucible melt 
about 25 grams of roll sulphur. Let it cool slowly, and when a 
thin crust has formed on the surface make a hole through this 
and pour out the liquid part of the sulphur. What is left ? 
Compare with the crystals formed in the last experiment. Lay 
the crucible aside, and in the course of a few days again examine 
the crystals. What changes, if any, have taken place ? 

Experiment 97. — Add hydrochloric acid to a solution of 
sodium thiosulphate. What takes place ? 

Experiment 98.— In a wide test-tube heat some sulphur to 
boiling. Introduce into it small pieces of copper-foil or sheet 
copper. Or hold a narrow piece of sheet copper so that the end 
just dips into the boiling sulphur. 

Experiment 99. — Dissolve some sulphur in concentrated caus- 
tic soda. In what form is the sulphur in the solution ? 



Hydrogen Sulphide. 

Experiment 100.— Arrange an apparatus as shown in Fig. 41. 
Put a small handful of the sulphide of iron, FeS, in the flask, 
and pour dilute sulphuric acid upon it. Pass the evolved gas 
through a little water contained in the wash-cylinder A. Pass 
some of the gas into water. [What evidence have you that it 
dissolves?] Collect some by displacement of air. Its specific 
gravity is 1.178. Set fire to some of the gas contained in a cyl- 
inder. In this case the air has not free access to the gas, and 
the combustion is not complete. The hydrogen burns to form 
water, while a part of the sulphur is deposited upon the inside 
walls of the cylinder. If there is free access of air, the sulphur 
burns to sulphur dioxide and the hydrogen to water. 

Make a solution of the gas in water in the usual way. Put 
some of this in a bottle and set it aside, and in the course of a 
few days examine it again. Boil another portion for a time in a 
test-tube, and note the changes. Pass a little of the gas through 
concentrated sulphuric acid contained in a test-tube, and note the 



HYDROGEN SULPHIDE: EXPERIMENTS. 



221 



changes. Moisten strips of paper with dilute solutions of lead 
nitrate, copper sulphate, stannous chloride, antimony chloride, 




Fig. 41. 



and mercuric chloride ; and expose these papers in turn to the 
gas. What changes take place ? Kepeat Experiment 90, and see 
whether one of the gases given off produces similar changes. 

Experiment 101. — Pass hydrogen sulphide successively through 
solutions containing a little lead nitrate, cadmium nitrate, and 
arsenic prepared by dissolving a little white arsenic, or arsenic 
trioxide, As 2 O s , in dilute hydrochloric acid. What action takes 
place in each case? The formula of lead nitrate is Pb(NO s )2; 
that of cadium nitrate, Cd(N0 3 ) 3 ; and that of the chloride of 
arsenic in solution is AsCl 3 . The corresponding sulphides are 
represented by the formulas PbS, CdS, and As a S 3 . 



CHAPTER XIV. 

COMPOUNDS OF SULPHUR, SELENIUM, AND TEL- 
LURIUM WITH OXYGEN AND WITH OXYGEN AND 
HYDROGEN. 

Introductory. — It has already been stated that, when 
the three elements of the sulphur group are burned in the 
air, they are converted into the corresponding dioxides. 
Under certain conditions it is possible to convert sulphur 
dioxide and tellurium dioxide into the trioxides S0 3 and 
Te0 3 , while the corresponding compound of selenium has 
not been made. A lower oxide of sulphur, S 2 3 , and a 
higher one of the formula S 2 7 have also been made. The 
dioxides dissolve in water, and from these solutions salts 
of the general formula M 2 S0 3 ,M 2 Se0 3 , and M 2 Te0 3 are 
obtained. The trioxide of sulphur dissolves in water with 
great evolution of heat, forming compounds S(OH) 6 , 
OS(OH) 4 , and0 2 S(OH) 2 . 

Most of the salts obtained from this solution are derived 
from an acid of the formula H SO, , and therefore this is 
generally called sulphuric acid. By treating sulphuric 
acid with reducing agents of various kinds it can be con- 
verted successively into other acids containing a smaller 
proportion of oxygen; aud if the reduction is pushed far 
enough sulphur and hydrogen sulphide are obtained. The 
limit of reduction is reached in hydrogen sulphide; and, 
on the other hand, if hydrogen sulphide is oxidized the 
limit of oxidation is reached in sulphuric acid. 

Sulphuric Acid, H 2 S0 4 . — Almost all of the salts of sul- 
phuric acid can be best explained on the assumption that 
the forms of sulphuric acid are derived from a compound 

222 






SULPHURIC ACID. 223 

S(OH) 6 . This is called normal sulphuric acid. From 
normal sulphuric acid by loss of water, the compounds 
H 4 S0 5 , H 2 S0 4 , and S0 3 are formed : 

S(OH) 6 =OS(OH) 4 +H 2 0; 
OS(OH) 4 = 2 S(OH), + H,0; 

2 S(OH) 2 = 3 S + H 2 0. 

While these compounds are all known, the salts of sul- 
phuric acid are for the most part derived from the acid 
containing two hydrogen atoms, viz., 2 S(OH) 2 . In some 
salts two of the hydrogen atoms in normal sulphuric acid 
are replaced by metals, the others remaining. Salts of the 
general formula S(OH) 4 (OM) 2 are thus formed. These are 
generally represented as containing two molecules of water 
of crystallization, thus, M 2 S0 4 -}- 2H 2 0. To decide be- 
tween the two views is at present impossible. In the first 
place, the question as to the nature of water of crystalliza- 
tion must be answered before it can be said whether there 
is any conflict between the two views. 

For the science as well as for the art of chemistry sul- 
phuric acid is of fundamental importance. It is used 
daily in every chemical laboratory and in every chemical 
factory, and in some of the most important branches of 
chemical industry enormous quantities of it are used. In 
consequence of the large demand for the acid the process 
used in preparing it has been studied with great care, and 
it has reached a high state of perfection. As it furnishes 
an excellent example of the applications of the facts of 
science to the building up of an industry, it will be studied 
with some degree of fulness. 

Sulphuric acid has been known for a long time. It was 
made in the eighteenth century by heatiug calcined iron 
vitriol (ferrous sulphate, FeS0 4 ) with sand, and was hence 
called oil of vitriol, a name which still survives. It was 
also prepared by treating sulphur with saltpetre (potassium 
nitrate, KN0 3 ). The acid was first prepared on the large 
scale in England with the use of comparatively large cham- 



224 COLLEGE CHEMISTRY. 

bers lined with lead. It was known as English sulphuric 
acid, and is still called by this name. 

Sulphuric acid occurs in nature in the form of salts or 
sulphates, such as calcium sulphate or gypsum, barium 
sulphate or heavy spar, and others. Although these salts 
occur in large quantity, the acid is not obtained from 
them, as there is no economical way of substituting 
hydrogen for the metal. The preparation of the acid by 
the substitution of hydrogen for the calcium in calcium 
sulphate, CaS0 4 , suggests itself, but this cannot be easi]y 
effected by any substance available in quantity. Hydro- 
chloric and nitric acids are made from their salts which 
occur in nature, the former, as we have seen, from sodium 
chloride, NaCl, the latter from saltpetre or potassium 
nitrate, KN0 3 , by treating with sulphuric acid. There 
is, however, no acid that acts upon the sulphates as sul- 
phuric acid acts upon chlorides, nitrates, and many other 
salts. 

The principal process that has up to recently been 
employed in the manufacture of sulphuric acid is based 
upon the two fundamental facts that (1) when sulphur is 
burned it is converted into sulphur dioxide, S0 2 ; and (2) 
when sulphur dioxide is treated with an oxidizing agent in 
the presence of water it is converted into sulphuric acid : 



so., 


+ 


H,0 


+ 





= H 2 SO,; 


or 


so, 


+ 


H.0 






= H 2 SO s ; 


an 


H 2 SO s 


+ 









= H 2 S0 4 . 





The chief difficulty is, of course, experienced in effecting 
the oxidation of the sulphurous acid. It is accomplished 
by introducing the gas, sulphur dioxide, into large cham- 
bers together with compounds of nitrogen and oxygen, 
and steam. That which plays the chief part in effecting 
the transformation is a mixture of the two oxides NO and 
N0 2 . This acts like the trioxide, N 2 3 , and it may be 
represented by this formula. Instead of starting with the 
trioxide, nitric acid is used in the manufacture of sul- 



SULPHURIC ACID. 225 

phuric acid. This at first reacts with sutyhur dioxide and 
steam, as represented in the equation 

2HNO, + 2S0 2 + H 2 = 2H 2 S0 4 + N 2 3 . 

After this the main reactions are (1) the formation of a 
compound of the formula S0 2 <^vr 2 , called nitrosyl- sul- 
phuric acid; and (2) the decomposition of the nitrosyl- 
sulphuric acid by water. These reactions are represented 
in the two equations following : 

2S0 2 + N 2 3 + 2 + H 2 = 2S0 2 (OH)(N0 2 ); 
2S0 2 < g£ + H 2 = 2S0 2 < gg + N 2 3 . 

The nitrogen trioxide formed in the second reaction again 
enters into combination with sulphur dioxide, oxygen, 
and water to form nitrosyl-sulphuric acid, which again 
undergoes decomposition with water. It will be seen, 
therefore, that the trioxide is not lost, but that it serves 
the purpose of effecting the oxidation of the sulphur 
dioxide, and, theoretically, a small quantity of the gas 
should be capable of transforming an infinite quantity of 
sulphur dioxide into sulphuric acid. 

Other reactions besides those mentioned above are 
undoubtedly involved in the manufacture of sulphuric 
acid, but, according to the most elaborate researches made 
on the subject in a factory in operation, those mentioned 
are the principal ones. Whatever the details may be, the 
oxidation is effected without difficulty, and the waste in 
nitrogen compounds is now but slight. 

The reactions, as has been said, are carried on in large 
chambers lined with lead, and known as the leaden cham- 
bers. The sulphur is burned in a special furnace so 
arranged that air has free access to it. The dioxide thus 
formed is then conducted through a tower so constructed 
that it presents a large surface to the action of the gas. 



2 2 6 



COLLEGE CHEMISTRY. 



Through this tower dilute sulphuric acid, taken from 
another tower at the end of the system, flows from above, 
and the hot gases coming in contact with this serve the 




purpose of concentrating it, while the gas itself is cooled 
before entering the leaden chamber. The general arrange- 
ment of the essential parts of a sulphuric -acid factory are 
shown in Fig. 42. 



SULPHURIC ACID. 227 

The sulphur dioxide passes through the large tube x 
into the tower G, called the Glover tower. This is filled 
from d to e with pieces of fire-brick over which from the 
cistern b a continual stream of dilute sulphuric acid flows. 
The gases are thus cooled down and the acid concentrated. 
From a second cistern there flows at the same time con- 
centrated sulphuric acid from the tower 67', or the Gay- 
Lussac tower. This stronger acid contains oxides of 
nitrogen in combination, and by contact with the dilute 
acid the oxides are set free and are thus mixed with the 
sulphur dioxide. The nitric acid is introduced into the 
first chamber, No. 1, in the form of vapor, together with 
the oxides of nitrogen and sulphur dioxide, and here also 
the gases meet with water-vapor. The reactions above 
referred to now take place. From the first chamber the 
gases pass through the pipe v into the second, from this 
into the third at w, and finally, in order to prevent the 
escape of any unused oxides of nitrogen, the gases are 
passed through the Gay-Lussac tower 67 '. This contains 
coke over which is kept flowing concentrated sulphuric 
acid, which takes up the oxides of nitrogen and will give 
them up again when diluted. This liberation of the oxides 
of nitrogen is accomplished, as has been said, in the 
Glover tower. The concentrated acid collected at the 
bottom of the Glover tower is well adapted for use in the 
Gay-Lussac tower. The leaden chambers in some factories 
are nearly 100 feet long, 18 to 30 feet broad, and 15 feet 
high. 

The acid taken from the chambers contains about 64 
per cent of the compound H 2 S0 4 , and has the specific 
gravity 1.5. This is~ evaporated first in lead pans until it 
reaches the specific gravity 1.75. As stronger acid acts 
upon lead, the evaporation beyond this point is carried on 
in platinum, glass, or iron. The strong acid thus obtained, 
which has a specific gravity of about 1.830, is the concen- 
trated sulphuric acid of commerce. 

Instead of sulphur, iron pyrites is now extensively used 
for the preparation of sulphur dioxide in the manufacture 



228 COLLEGE CHEMISTRY. 

of sulphuric acid. This is a compound of iron and sulphur 
of the composition FeS 2 . When it is heated in contact 
with the air it is converted into the oxide, and the sulphur 
passes off in the form of sulphur dioxide. If the sulphide 
used contains selenides the selenium dioxide formed in the 
roasting process is carried into the flues, and is there 
deposited with the flue-dust, as was stated in speaking of 
the source of selenium. 

Commercial sulphuric acid prepared in the above way 
is an oily liquid, usually somewhat colored by impurities. 
The nature of the impurities is depeudent upon the sub- 
stances used in the manufacture and upon the conditions. 
It often contains some lead sulphate in solution, and when 
it is diluted with water this separates, giving the liquid a 
more or less cloudy appearance. By standing, however, 
the liquid becomes clear, as the lead sulphate .collects at 
the bottom. For obvious reasons, some oxides of nitrogen 
are also generally present. Among other impurities fre- 
quently met with in the commercial acid are arsenic from 
the pyrites, and a little selenium. 

A simpler process than that above described has recently 
been introduced. This promises gradually to displace the 
older method, as it has already been adopted by a number 
of large factories. It consists in passing sulphur dioxide 
and air together over finely-divided platinum, ferric oxide, 
or some other substance that has the power to effect the 
union of the sulphur dioxide and the oxygen of the air. 
The direct product of the union is sulphur trioxide, S0 3 . 
This with water gives sulphuric acid. The reactions in- 
volved are : 

SO, + = SO s ; 
SO~ + H 2 = H 2 S0 4 . 

Pure Sulphuric Acid is made by the method just 
described or from the commercial acid by treating it with 
such substances as will remove the oxides of nitrogen and 
arsenic, and by distilling. It is a colorless liquid at the 
ordinary temperature. The pure acid generally made has 



SULPHURIC ACID. 229 

about the same concentration as the commercial crude 
arid. By taking special precautions in the distillation a 
product having very nearly the composition H 8 SO l can be 

obtained. This is a thick, clear liquid of specific gravity 
1.854 at 0°. When cooled down to a low temperature it 
forms large crystals which melt at 10.5°. When heated 
it gives off some sulphur trioxide in consequence of partial 
decomposition into this substance and water: 

H 2 S0 4 = H 2 + S0 3 . 

It finally boils, however, at the temperature 338°, and the 
distillate has the composition represented by the formula 
H 2 S0 4 -j- y^E^O. Heated somewhat above its boiling- 
point it is completely decomposed into sulphur trioxide 
and water, according to the above equation. If the heat- 
ing is carried to a higher temperature the sulphur trioxide 
breaks down into sulphur dioxide and oxygen. 

Sulphuric acid has a strong tendency to absorb water, 
and to form compounds with it. In consequence of the 
formation of these compounds a great deal of heat is 
evolved when sulphuric acid is mixed with water. One 
molecule of sulphuric acid when mixed with about 1600 
molecules of water gives 17,850 cal. Among the com- 
pounds thus formed are the so-called hydrates of the com- 
position H 2 S0 4 -f H 2 0, and H 2 S0 4 + 2H 2 0, which should 
probably be regarded as having the constitution OS(OH) 4 
and S(OH) 6 . 

So strong is the tendency of the acid to combine with 
the elements of water that it abstracts them in the propor- 
tions to form water from many organic substances. Thus 
a piece of wood placed in sulphuric acid soon turns black 
and is completely disintegrated. The cause of this is to 
be found in the fact that wood consists essentially of 
carbon, hydrogen, and oxygen; and the sulphuric acid 
abstracts the hydrogen and oxygen, leaving the carbon 
mainly in the uncombined state. Similarly it abstracts 
moisture from gases, and it is used extensively in the 



230 COLLEGE CHEMISTRY. 

laboratory for this purpose. In contact with the skin it 
acts as it does upon wood, causing wounds which are pain- 
ful and difficult to heal. 

Sulphuric acid is called a strong acid, a term which 
needs further explanation, and the subject of the relative 
strengths of acids will be discussed in due time. As used 
here, it means simply that the acid has the power to 
decompose the salts of most other acids, appropriating the 
metals and setting the other acids free. This is illustrated 
by the formation of hydrochloric and nitric acids by treat- 
ing sodium chloride and potassium nitrate respectively with 
sulphuric acid. We shall see, however, that the fact that 
sulphuric acid decomposes the salts of hydrochloric and 
nitric acids does not prove that it is a stronger acid than 
they are. There are other facts besides the strengths of 
the acids which determine whether such decompositions 
take place or not. 

Sulphuric acid gives up its oxygen to other substances 
comparatively easily, and is generally reduced to sul- 
phurous acid, which is decomposed into sulphur dioxide 
and water. Thus hydrogen sulphide and hydrobromic 
acid both act upon it, as we have seen; the products 
being, in the former case, sulphur dioxide, water, and 
sulphur; in the latter, sulphur dioxide, water, and bromine: 

H 2 S0 4 + H 2 S = 2H 2 + S0 2 + S; 
H 2 S0 4 + 2HBr = 2H 2 + S0 2 + Br 2 . 

Generally, in acting upon metals it forms the correspond- 
ing salt and hydrogen is given oif, but if the temperature 
is high and the acid concentrated the hydrogen acts upon 
the acid, reducing it. Thus, in making hydrogen by 
treating zinc with sulphuric acid, if the acid is dilute and 
is kept cool the reaction takes place in the simplest way; 
but if the acid is concentrated and is allowed to get hot 
some hydrogen sulphide is always formed. Copper doss 
not act upon sulphuric acid at ordinary temperatures. If 
it is treated with concentrated sulphuric acid, tulphur 



TETRAHYDROXYL-SULPHURIC ACID. 231 

dioxide is the chief reduction-product. Even free hydrogen 
if passed into sulphuric acid heated to 160° reduces it, 
forming sulphur dioxide. The complete reduction of sul- 
phuric acid to sulphur and hydrogen sulphide is beautifully 
shown by the action of hydriodic acid. As was stated in 
speaking f the action of sulphuric acid upon potassium 
iodide, four reactions may take place when these sub- 
stances act upon each other. They are represented by the 
equations 

2KI + H 2 S0 4 = K 2 S0 4 + 2III; 
H 2 S0 4 + 2111 = 21I.,0 + S0 2 + 21; 
H 2 S0 4 + GUI = 411,0 + S +61; 
H 2 S0 4 + 8HI = 41I 2 + H 2 S f 81. 

Carbon and sulphur also act upon sulphuric acid and 
reduce it to sulphur dioxide. With sulphur the action is 
represented thus: 

2H 2 S0 4 + S = 3S0 2 + 2H 2 0. 

Sulphuric acid is a dibasic acid (see p. 162) and forms 
two series of salts, normal salts of the general formula 
M 9 S0,, and acid salts of the formula MHSO.. The sul- 
phates are very stable salts. Those of the strongest bases 
are not decomposed by the highest heat. Those of weaker 
bases break down, giving up sulphur trioxide ; or if the 
decomposition takes place at a temperature above that at 
which sulphur trioxide breaks down, this is decomposed 
into sulphur dioxide and oxygen. 

Tetrahydroxyl-Sulphuric Acid, 0S(0H) 4 .— This com- 
pound has already been referred to as being formed by the 
addition of water to ordinary sulphuric acid. It is a 
crystallized compound which melts at 7.5°. No salts of 
this acid are known, or, rather, no salts derived from it 
by the replacement of all the hydrogen by metal are 
known. Only two of the hydrogen atoms appear to be 
replaceable. This compound is generally represented by 
the formula II 2 S0 4 + H 2 0, and called a hydrate. 



232 COLLEGE CHEMISTRY. 

Normal Sulphuric Acid, S(OH) 6 . — This is commonly 
represented by the formula H 2 S0 4 -f- 2H 2 and, like the 
preceding compound, is regarded as a hydrate. It appears 
to be formed by the action of water on sulphuric acid. 
On mixing sulphuric acid and water the maximum con- 
traction takes place when the quantities necessary to form 
this compound are brought together, and there are other 
good reasons for believing that the compound exists in the 
solution. It is not a solid like the preceding compound. 
The acid forms no salts bearing simple relations to it. 

Disulphuric Acid, Pyrosulphuric Acid, H 2 S 2 7 . — This 
compound, which is also known by the names fuming 
sulphuric acid and Nordhausen sulphuric acid, is closely 
related to ordinary sulphuric acid, and is made from it by 
treating it with sulphur trioxide, the two combining 
directly, as represented thus: 

H 2 S0 4 + S0 3 = H 2 S 2 7 . 

It is made by distilling ferrous sulphate which is not per- 
fectly dry : 

4FeS0 4 + li,0 = 2Fe 2 3 + 2S0 2 + H,S,O r 

A so-called solid sulphuric acid is now manufactured by 
a process which will be referred to under Sulphur Trioxide. 
It is essentially disulphuric acid. 

Disulphuric acid, as it is found in the market, is gen- 
erally a thick liquid which gives off dense fumes when 
exposed to the air, and breaks down completely into sul- 
phur trioxide and sulphuric acid when heated. When 
pure it crystallizes in large crystals which melt at 35°. 

It is believed that the relation between disulphuric acid 
and ordinary sulphuric acid should be expressed by these 
formulas : 



0£<oh+HO >80 » = ' 8 \ 

uii av \ 0H H0 . 




SULPHUROUS ACID. 233 

Or the formula of the acid may be written thus: 

2 S-OH 

1 


2 S-OH. 

Sulphurous Acid, H 2 SO s . — While no acid of the formula 
H 2 S0 3 is known in the free condition, a large number of 
salts which are derived from this acid are known. They 
are made by treating a water solution of sulphur dioxide 
with bases, and therefore it is believed that the solution 
contains the acid which is formed by the action of sulphur 
dioxide on water, thus: 

S0 2 + H 2 = H 2 S0 3 . 

It is, however, so unstable that it breaks down into the 
dioxide and water at every attempt to isolate it. The 
dioxide, as has been stated and as will be shown more fully, 
is formed by the burning of sulphur and by the reduction 
of sulphuric acid. The acid forms a number of unstable 
hydrates, apparently of complicated composition. Owing 
to their great instability, however, the investigation of 
these substances is extremely difficult and unsatisfactory. 

Sulphurous acid takes up oxygen readily and is thus 
transformed into sulphuric acid. It is only necessary to 
allow a solution to stand for a time to find that the odor 
of the gas disappears and that sulphuric acid is then present 
in the solution. Sulphurous acid is frequently used in the 
laboratory as a reducing agent. 

The sulphites take up oxygen to form sulphates, and 
they also take up sulphur to form thiosulphates. The 
two reactions appear to be perfectly analogous: 

Na 2 S0 3 + =-. Na 2 S0 4 ; 

Na 2 S0 3 + S = Na 2 S 2 3 (or N 2 S0 3 S). 

Sulphurous acid forms two series of salts, the normal 
sulphites of the general formula M 2 S0 3 , and the acid sul- 



234 COLLEGE CHEMISTRY. 

phites of the general formula MHS0 3 . These are unstable, 
though much more stable than the acid itself. When 
treated with most acids they are decomposed, yielding 
sulphur dioxide instead of sulphurous acid. The decom- 
position of sodium sulphite with hydrochloric acid is 
represented by the equation 

Na 2 S0 3 '+ 2HC1 = 2NaCl + H 2 + S0 2 ; 

with sulphuric acid thus: 

Na 2 S0 3 + H 2 S0 4 = Na 2 S0 4 + H,0 + S0 2 . 

Hyposulphurous Acid, H 2 S 2 4 . — This acid is also called 
hydrosulphurous acid, but the name hyposulphurous acid 
is more in accordance with the nomenclature adopted for 
the other acids, and is now preferred. But little is known 
of the compound. It is formed by reduction of a salt of 
sulphurous acid by zinc: 

Zn + 2S0 2 = ZnS 2 4 . 

The free acid cannot be obtained from this salt. 

Thiosulphuric Acid, H 2 S 2 3 . — This acid was formerly, 
and is still by many, called hyposulphurous acid. Its 
formation, or the formation of its salts by the addition of 
sulphur to the sulphites, has been mentioned, and the 
analogy between this reaction and that of the formation of 
sulphates by the addition of oxygen to sulphites has been 
commented upon. It may be regarded as sulphuric acid 
in which one atom of sulphur has been substituted for one 
atom of oxygen, and hence the name thiosulphuric acid 
is appropriate, whereas the name hyposulphurous acid 
suggests at once a compound similar to' sulphurous acid, 
but containing less oxygen, and is therefore inappropriate. 

The acid itself is very unstable, breaking down into 
sulphur dioxide, sulphur, and water. By acids its salts 
are decomposed in a similar way with evolution of sulphur 
dioxide and separation of sulphur, which appears sus- 



SULPHUR DIOXIDE. 235 

pended in the liquid in ;i very fine state of division. With 
hydrochloric acid the decomposition takes place thus: 

Na a S 2 O s + 2HC1 = 2NaCl + SO, 4- S + H 2 0. 

Other Acids of Sulphur. — Of the other acids of sulphur 
but little, need be said here. They form a series the 
members of which are closely related to one another as 
shown below: 

Dithionic acid H 2 S 2 6 

Trithionic acid H 2 S 3 6 

Tetrathionic acid H 2 S 4 6 

Pentathionic acid H 2 S 5 G 

Persulphuric Acid, HS0 4 , is formed by dissolving the 
oxide S 2 7 in water, the oxide itself being formed by sub- 
jecting a mixture of sulphur dioxide and oxygen to the 
silent discharge in an ozone tube (see p. 11G). The potas- 
sium salt of persulphuric acid is easily obtained by sub- 
jecting a saturated solution of acid potassium sulphate to 
electrolysis. 

Caro's Reagent. — When potassium persulphate is added 
to moderately dilute sulphuric acid a solution is obtained 
that shows remarkable oxidizing powers. It has recently 
come into use in laboratories under the name of Carol's 
reagent. 

Compounds of Sulphur with Oxygen. — Sulphur, as has 
been repeatedly stated, combines with oxygen in two pro- 
portions, forming the oxides S0 2 and S0 3 , or sulphur 
dioxide and sulphur trioxide. Besides these two it also 
forms a sesquioxide, S 2 3 , and a heptoxide, S 2 7 , but com- 
paratively little is known in regard to the last two. The 
one best known is sulphur dioxide. 

Sulphur Dioxide, S0 2 . — This, as has been seen, is 
formed when sulphur is burned in the air or in oxygen; 
and it is also easily formed by reduction of the higher 
oxides and acids of sulphur. Owiug to the fact that with 



236 COLLEGE CHEMISTRY. 

water it forms sulphurous acid, it is frequently called sul- 
phurous anhydride. The methods for making it were 
referred to under sulphuric acid, and the reactions involved 
were sufficiently discussed. It need only be said here that 
in the laboratory the methods most commonly employed 
are: (1) Heating sulphuric acid with copper: (2) heating 
the acid with carbon (charcoal); (3) heating the acid with 
sulphur; and (4) allowing concentrated sulphuric acid to 
drop into a 40 per cent solution of acid sodium sulphite : 

H 2 S0 4 + HNaS0 3 = HNaSO; + S0 2 + H 2 0. . 

Sulphur dioxide is a colorless, transparent gas, which 
has a pungent, suffocating odor, familiar as the odor of 
burning sulphur matches. It is poisonous, causing death 
when inhaled in any quantity, and giving rise to bad 
symptoms even in comparatively small quantities. It does 
not readily give up its oxygen, so that burning bodies are 
extinguished when introduced into it. It acts something 
like water in this respect. It is more than twice as heavy 
as air, its specific gravity being 2.26. When sulphur is 
burned in oxygen gas the sulphur dioxide formed occupies 
the same volume as the oxygen used up, so that there is 
no change in the volume. This can be shown by the 
experiment here described. In a bent 
glass tube, of the form shown in Fig. 
43, there is placed a piece of sulphur, 
and the tube is then half filled with 
pure oxygen over mercury. On now Fro. 43. 

heating the tube at the part where the sulphur is, this 
burns and is converted into sulphur dioxide. After the 
tube has cooled down to the ordinary temperature the gas 
is found to occupy the same volume as before. This will 
be readily understood by the aid of the following con- 
siderations : In the reaction 

s + o 2 = so 2 

one molecule of sulphur dioxide is formed for every mole- 




SULPHUR DIOXIDE. 237 

cule of oxygen used up. But a molecule of sulphur dioxide 
in the form of gas occupies the same volume as a molecule 
of oxygen, so that, as the volume occupied by the sulphur 
in the experiment is insignificant, there is no change in 
volume occasioned by the above reaction. 

Sulphur dioxide dissolves in water, as we have seen, and 
forms a liquid in which, judging by its conduct, sul- 
phurous acid is present. 

The gas is easily liquefied by cold alone. It is only 
necessary for this purpose to pass the dry gas through a 
tube surrounded by a freezing mixture of ice and salt. 
The liquid changes rapidly into gas under ordinary pressure 
at the ordinary temperature. In this change so much heat 
is absorbed that a temperature of about — 60° can be pro- 
duced by means of it; and a portion of the liquid can be 
solidified. 

Sulphur dioxide is very stable. If heated to 1200° 
under pressure, however, it breaks down into sulphur tri- 
oxide and sulphur: 

3S0 2 = 2S0 3 + S. 

When conducted into solutions of bases or of carbonates 
the corresponding sulphites are formed: 

2KOH + SO, = K 2 S0 3 + H 9 0; 
K 2 C0 3 + S0 2 = K 2 S0 3 + C0 2 . 

Under certain conditions, as when the gas is passed into 
a hot solution of an alkali carbonate, a salt of the general 
formula M 2 S 2 0. is formed. This bears to the sulphite the 
same relation that the pyrosulphate bears to the sulphate : 

2KHSO3 - K 2 S 2 5 + H 2 0; 

and 2KHSO, = K 2 S 2 7 + H 2 0. 

Sulphur dioxide is used extensively for the purpose of 
bleaching silk, wool, straw, and basket-ware. In order 
that it may bleach, however, water must be present, so 



238 COLLEGE CHEMISTRY. 

that it appears that the true bleaching agent in this case 
is sulphurous acid and not the dioxide. When we consider 
that sulphur dioxide does not readily take up nor give up 
oxygen, while sulphurous acid does readily take it up, the 
necessity of having water present in the bleaching process 
at once becomes apparent. The bleaching in some cases 
certainly consists in abstracting oxygen from the colored 
substances, and thus converting them into colorless 
products. In other cases it is due to the formation of 
compounds of sulphurous acid with the d3^e-stuffs. 

Sulphur dioxide is not only a bleaching agent like 
chlorine, but like chlorine it is also a disinfectant. It has 
to some extent the power to destroy the organisms which 
cause changes in organic substances. It prevents fermen- 
tation and is therefore used as a preservative. Its power 
to destroy the germs of disease, that is, to disinfect, is not 
as great as is frequently supposed. Much larger quanti- 
ties are necessary for this purpose than are commmonly 
used. 

Sulphur Trioxide, S0 3 . — This compound is made by 
passing sulphur dioxide and oxygen together over heated 
platinum in a finely-divided state. It is obtained most 
readily by heating disulphuric acid, which breaks up easily 
into sulphur trioxide and ordinary sulphuric acid according 
to the equation 

H 2 S 2 7 = H 2 S0 4 + S0 3 . 

Similarly, the acid sulphates of the alkali metals yield the 
corresponding normal sulphates and sulphur trioxide : 

2NaHS0 4 = Na 2 SO, + S0 3 + H 2 0. 

It is now manufactured on the large scale by passing sul- 
phur dioxide and oxygen together over asbestos covered 
with finely-divided platinum, ferric oxide, and some other 
substances, and the product thus obtained is converted into 
sulphuric acid by the action of water, or it is passed into 



SULPHUR TRIOXIDE. 239 

ordinary sulphuric acid for the purpose of making "solid 
sulphuric acid," which is almost pure disulphuric acid, 

Sulphur trioxide is a white crystallized solid which 
appears to exist in two modifications. The oue crystallizes 
in line needles like asbestos. When heated it passes 
directly into gas. The other solid variety melts at 14.8°, 
forming a liquid which boils at 46°. In contact with 
the air the oxide gives off thick fumes which are partly 
due to the great power of the compound to combine with 
water. Water acts with violence upon it, the heat evolved 
in the act being 39,170 cal. It also acts upon substances 
containing hydrogen and oxygen in much the same way 
that concentrated sulphuric acid does, charring them by 
abstracting the hydrogen and oxygen. It acts, however, 
more violently in this way than sulphuric acid does. 
With water it forms sulphuric acid, and it is, therefore, 
called sulphuric anhydride. The reaction involved in 
passing from sulphur trioxide to sulphuric acid is of a 
kind which, as we have seen, is frequently met with both 
with acidic oxides or anhydrides, and with basic or metal- 
lic oxides; and it is desirable that it should here be 
studied a little more carefully than it has yet been. What 
we know is that when sulphur trioxide acts upon water 
there is a great deal of heat evolved, and compounds of 
different composition are obtained. The composition of 
these compounds is represented by the formulas S0 3 -|- 
H 2 0, S0 3 -f 2H 2 0, and S0 3 -f 3H 2 0. So, too, when 
calcium oxide or lime, CaO, acts upon water, there is great 
evolution of heat, and a compound is formed the composi- 
tion of which is represented by the formula CaO + H 2 0. 
But in these formulas no attempt is made to give any 
account of what takes place in the chemical acts referred 
to. That the water is not present in the compounds as 
water seems evident, in the first place from the conduct of 
the compounds, and in the second place from the amount 
of heat evolved in the act of combination. Now, taking 
the chemical conduct of the substances into consideration, 



240 COLLEGE CHEMISTRY. 

they appear to contain hydrogen in combination with 
oxygen, and their conduct becomes comprehensible on the 
supposition that the group known as hydroxyl (-0-H) is 
present. This view has been found to be in accordance 
with a large number of facts, and it is of great assistance 
in dealing with these facts. The view is distinctly this: 
When an acidic oxide acts upon water it is converted into 
a hydroxyl compound which has acid properties, as shown 
m this equation: 





II H 2 

= S = -f H 2 



H 2 



f OH 
OH 
OH 
OH' 
OH 
OH 



According to this, each molecule of Avater is decomposed 
and each hydrogen atom in the resulting compound is in 
combination with oxygen. The same kind of action takes 
place in the case of some basic oxides, as shown, for ex- 
ample, in the case of calcium oxide : 

Ca=0 + H 3 = Ca<^|. 

Just as sulphur trioxide acts upon water to form sul- 
phuric acid, so it acts upon metallic oxides, to form 
sulphates. Thus with calcium oxide it forms calcium 
sulphate: 

CaO + S0 3 = CaS0 4 . 

Sulphuryl Chloride, S0 2 C1 2 , is formed by the direct 
action of chlorine upon sulphur dioxide in the sunlight, 
the action being similar to that which takes place when 
sulphur dioxide and oxygen unite to form the trioxide: 

so, + a = S0 2 C1 2 ; 
S0 o + = SO,. 



COMPOUNDS OF SELENIUM AND TELLURIUM. 241 

It is best prepared by heating the compound S0 2 <}di 
to 170-180 , when the following reaction takes place: 

2S0 2 <g 1 H = H a S0 4 + SO a Cl 2 . 

It is a liquid which is easily decomposed by water, as 
represented in the equation 

S0 2 C1 2 + 2H 2 = S0 2 (OH) 2 + 2HC1. 

With half the quantity of water required to effect the 

PI 
above decomposition chlorsulphuric acid, S0 2 <^ H , is 

formed : 

S0 2 <£{ + H 2 = S0 2 <^ H + HC1. 

Chlorsulphuric Acid, or Sulphuryl-hydroxyl Chloride. 
S0 2 <qtj, is formed by the direct action of hydrochloric 
acid upon sulphur trioxide : 

SO3 + HC1 = so 2 <£ 1 H . 

Like sulphuryl chloride it is decomposed by water, yielding 
hydrochloric acid and sulphuric acid. 

Compounds of Selenium and Tellurium with Oxygen 
and with Oxygen and Hydrogen. — In the introduction to 
this chapter it was stated that selenium and tellurium form 
compounds with oxygen corresponding to sulphur dioxide 
and sulphur trioxide. Both of these oxides of tellurium 
are known, together with a third of the composition repre- 
sented by the formula, TeO, while only one oxide of 
selenium, the dioxide. SeO., . is known. On the other 
hand, the acids of selenium and tellurium, corresponding 
to sulphurous and sulphuric acids, are known. 



242 COLLEGE CHEMISTRY. 

Selenious Acid, H 2 SeO s . — This compound is formed by 
oxidation of selenium in presence of water or by dissolving 
selenium dioxide in water. It is a crystallized solid, which 
attracts moisture from the air, and when heated breaks 
down into selenium dioxide and water. It forms two 
series of salts, the acid selenites, MHSe0 3 , and the normal 
selenites, M 2 Se0 3 . While in composition the acid and its 
salts are analogous to sulphurous acid and the sulphites, 
in conduct there is a marked difference. Sulphurous acid 
tends, as we have seen, to take up more oxygen and form 
sulphuric acid, while selenious acid gives up its oxygen 
with great ease and yields selenium. When a solution of 
selenious acid is treated with sulphur dioxide the acid is 
reduced to selenium : 

H 2 Se0 3 + 2S0 2 + H 2 = 2H 2 S0 4 + Se. 

Selenic Acid, H 2 Se0 4 , is formed by the action of power- 
ful oxidizing agents like saltpetre on selenium. Chlorine 
and bromine in water solution also serve the purpose. 
The action with chlorine is represented thus : 

Se + 601 + 4H 2 = H 2 Se0 4 + 6H01. 

Selenium Dioxide, Se0 2 . — This analogue of sulphur 
dioxide is made by burning selenium, or by heating sele- 
nious acid. It is a solid crystallized substance that can be 
sublimed without undergoing decomposition and without 
melting. It dissolves readily in water, forming selenious 
acid. When it is sublimed particles of dust which may be 
present in the vessel effect partial reduction, and the 
product, instead of being white, as the oxide is when pure, 
is colored by the particles of selenium. 

Tellurious Acid, H 2 Te0 3 . — Tellurious acid is formed by 
treating tellurium tetrachloride with water. It is possible 
that the first action causes the formation of normal tel- 
lurious acid, Te(OH) 4 , and that this then breaks down 
into tellurious acid, H 2 Te0 3 , and water: 



TELLURIC ACID. 243 

CI HOH f OH 

, CI . HOH r V J OH . AUm 
Te id + HOH = iM 0H + 4HCl; 

L Cl HOH [ OH 

Te ^ gg = Te OH + H,0. 
[OH ( U 

The potassium salt of the acid is formed by melting 
together tellurium dioxide and potassium carbonate : 

K 2 C0 3 + Te0 2 = K 2 TeO + C0 2 . 

If the salt thus formed is dissolved in water and nitric acid 
added to the solution, tellurious acid is thrown down : 

K 2 Te0 3 + 2HNO s = 2KN0 8 + H 2 Te0 3 . 

It is a solid which easily loses water and is thus trans- 
formed into tellurium dioxide. 

Telluric Acid, H 2 Te0 4 . — Salts of this acid are formed by 
melting tellurious acid with saltpetre and other oxidizing 
agents. When the solution of the acid is evaporated to 
crystallization the solid compound deposited has the com- 
position H 6 Te0 6 , and, according to what was learned in 
studying sulphuric acid, it appears probable that this is 
normal telluric acid, Te(OH) 6 . When normal telluric acid 
is heated to a little above 100° it loses water and is trans- 
formed into the acid H 2 Te0 4 , corresponding to ordinary 
sulphuric acid, from which most of the tellurates are 
derived : 

Te(OH) 6 = Te0 2 (OH) 2 + 2H 2 0. 

Heated higher, to about 160°, the acid is decomposed into 
tellurium trioxide and water: 

TeO a (OH) 2 = Te0 3 + H 2 0. 



244 COLLEGE CHEMISTRY. 

Although most of the tellurates are simple salts of the 
acid H 2 Te0 4 , others are derived from more complex forms 
of the acid. One of these is analogous to disulphuric acid. 

Oxides of Tellurium. — Tellurium monoxide is formed 
by heating- sulphur trioxide and tellurium together in a 
vacuum. Tellurium dioxide is formed by burning tel- 
lurium or by oxidizing it with nitric acid; and, further, 
by the decomposition of tellurious acid by heat. It 
crystallizes and is but slightly soluble in water. 

The trioxide, Te0 3 , is formed by heating telluric acid 
to a high temperature. Its conduct is entirely different 
from that of sulphur trioxide. While the latter acts with 
violence upon water, and readily upon metallic oxides and 
hydrochloric acid, the former does not act readily upon 
any of these substances. It is insoluble in hot as well as 
cold water. 



Family VI, Group A. 

Group A, Family VI, includes chromium, molybdenum, 
tungsten, and uranium. All of these show some resem- 
blance to the elements of the sulphur group, but they also 
appear in entirely different characters, forming compounds 
of a kind unknown among the derivatives of sulphur and 
its analogues. The relation which these elements bear to 
sulphur is much like that which manganese bears to 
chlorine. The resemblance to sulphur is seen mainly in 
the formation of acids of the formulas H 2 Cr0 4 , H 2 Mo0 4 , 
H 2 W0 4 , and H 2 U0 4 ; and the oxides Cr0 3 , Mo0 3 , WO s , 
and U0 3 . The common salts of chromic acid are derived 
from dichromic acid, which is analogous to disulphuric 
acid. They have the general formula M 2 Cr 2 7 . So, too, 
salts of molybdic acid are known which are derived from 
the simple form of the acid, H 2 Mo0 4 , and others which 
are derived from a dimolybdic acid, H 2 Mo 2 7 , and from 
more complicated forms. Tungsten has a wonderful 
power of forming complex acids. All of them, however, 



FAMILY VI, GROUP A. 245 

can be referred to the simple form H 2 W0 4 . And, finally, 
uranic acid forms suits which for the most part are derived 
from diuranic acid, H 2 U 2 O r All the most important of 
these compounds will be taken up later. 

When the acids of chromium, molybdenum, tungsten, 
and uranium lose oxygen they form compounds that have 
little or no acid character. The lower oxides of chromium 
form salts with acids, and these bear a general resemblance 
to the salts of aluminium, iron, and manganese. The 
eliro mates lose their oxygen quite readily when acids are 
present with which the chromium can enter into combina- 
tion in its capacity as a base-forming element. Thus, 
when potassium chromate, K 2 Cr0 4 , is treated w>\th hydro- 
chloric acid in the presence of something which can take 
up oxygen, decomposition takes place thus: 

2K 2 Cr0 4 + 10HC1 = 4KC1 + 2CrCl 3 + 5II 2 + 30. 

With sulphuric acid the action takes place as represented 
in this equation: 

2K 2 Cr0 4 + 5H 2 S0 4 = 2K 2 S0 4 + Cr 2 (SOJ 3 + 5H 2 + 30. 

In both these cases the chromium enters into combination 
as a trivalent base-forming element, taking the place of 
three atoms of hydrogen in hydrochloric acid in the first 
case, and of three atoms of hydrogen in the sulphuric acid 
in the second. Molybdenum and tungsten do not form 
salts of this character; indeed they seem to be practically 
devoid of basic properties. Uranium, on the other hand, 
forms some curious salts which differ from the simple 
metallic salts which we commonly have to deal with. 



246 



COLLEGE CHEMISTRY. 



EXPERIMENTS. 

Manufacture of Sulphuric Acid. 

Experiment 102.— The manufacture of sulphuric acid can be 
illustrated in the laboratory by means of the apparatus repre- 
sented in Fig. 44. This consists of a large balloon-flask fitted 
with a stopper having five openings. By means of tubes it is con- 
nected with three small flasks. One of these, a, contains water 
for the purpose of providing a current of steam ; another, c, con- 
tains copper-foil and concentrated sulphuric acid, which give sul- 




Fig. 44. 

phur dioxide when heated ; and the third, 6, contains copper- foil 
and dilute nitric acid, which give oxides of nitrogen, mainly 
nitric oxide, NO. When the nitric oxide comes in contact with 
the air it combines with oxygen, forming nitrogen trioxide and 
nitrogen peroxide ; and when steam and sulphur dioxide are ad- 
mitted to the flask the reactions involved in the manufacture of 
sulphuric acid take place. By means of a pair of bellows attached 
at d air is supplied. If air is not forced in, the gases become 
colorless, owing to complete reduction of the oxides of nitrogen 
to the form of nitric oxide, NO, which is colorless. If steam is 
not admitted the walls of the vessel become covered with crystals 
of nitrosyl-sulphuric acid. This is, however, decomposed by an 
excess of steam. 



SULPHUROUS ACID AND SULPHUR DIOXIDE. 247 

EXPERIMENT 103.— Into a vessel containing ordinary concen- 
trated sulphuric acid introduce small sticks of wood, pieces of 
paper, and various other organic substances, and note the result. 
The charring effect is particularly well shown by adding the acid 
drop by drop to a concentrated solution of sugar, or to molasses, 
and stirring. 

Experiment 104. — Sulphuric acid is detected in analysis by add- 
ing barium chloride to its solution, when insoluble barium sul- 
phate is formed: 

H 2 S0 4 + BaCl a = BaSO< + 2HC1. 

Other insoluble sulphates are those of strontium and lead ; and 
calcium sulphate is difficultly soluble. To a dilute solution of sul- 
phuric acid or of any soluble sulphate, add in test-tubes barium 
chloride, strontium nitrate, and lead nitrate. 

Sulphurous Acid axd Sulphur Dioxide. 

Experiment 105. — Put eight or ten pieces of sheet copper, one 
or two inches long and about half an inch wide, in a 500 cc. flask; 
pour 15 to 20 cc. concentrated sulphuric acid on it. On heating, 
sulphur dioxide will be evolved. The moment the gas begins to 
come off, lower the flame, and keep it at such a height that the 
evolution is regular and not too active. Pass some of the gas into 
a bottle containing water. The solution in water is called sul- 
phurous acid. 

Experiment 106. — The most convenient method for making 
sulphur dioxide consists in letting moderately concentrated sul- 
phuric acid drop into a water solution of acid sodium sulphite. 
Fit a flask with a stopper with two holes. Through one of the 
holes pass a dropping-funnel and through the other a convenient 
delivery-tube. In the funnel put "a cooled mixture of equal vol- 
umes of concentrated sulphuric acid and water. Tn the flask put 
a 40-per cent solution of acid sodium sulphite in water. By 
turning the stop-cock of the funnel so that the acid drops into the 
solution there will be a regular evolution of the gas. 

Experiment 107. — Pass sulphur dioxide into a moderately dilute 
solution of potassium hydroxide, until the solution is saturated. 
"What is then contained in the solution ? To a little of it add 
hydrochloric acid. What takes place ? 

Experiment 108.— Try the effect of heating concentrated sul- 
phuric acid with charcoal, and with sulphur. 



24S 



COLLEGE CHEMISTRY. 



Experiment 109.— Collect by displacement of air some of the 
gas made in Experiment 106. Does it burn ? or does it support 
combustion ? 

Experiment 110. — Pass some of the gas through a bent-glass 
tube surrounded by a freezing mixture of salt and ice. Tubes 




Ftg 45. 

provided with glass stop-cocks are made for such purposes. They 
generally have the form represented in Fig. 45. If the tube is 
taken out of the freezing mixture, the liquid sulphur dioxide 
changes rapidly to gas, if the tube is open. 

Experiment 111.— Burn a little sulphur in a porcelain crucible 
under a bell- jar u Place over the crucible on a tripod some flowers. 
In the atmosphere of sulphur dioxide the flowers will be bleached. 

Sulphurous Acid is a Reducing Agent. 

Experiment 112. — To a dilute solution of potassium iodide in 
a test-tube gradually add chlorine water until the solution be- 
comes clear and colorless. Now add a solution of sulphurous acid. 
At first iodine is deposited, but on further addition of sulphurous 
acid it dissolves again. Explain all the changes. 

Sulphur Trioxide. 



Experiment 113. — Heat a little fuming sulphuric acid gently in 
a test-tube. What takes place ? Put a little of the acid (5-10 cc.) 
in a small dry retort provided with a glass stopper and connect 
with a dry glass receiver. Heat the retort gently, and keep the 
receiver cool. By means of a dry glass rod take out some of the 
substance which collects in the receiver and put it in water. Lay 
a little of it on a piece of wood and on a piece of paper. 

Experiment 114. — Prepare finely-divided platinum by moisten- 
ing some fine asbestos with a solution of platinic chloride and 
heating to redness in a porcelain crucible. The substance thus 
obtained is known as platinized asbestos. Now arrange an 



SULPHUR T RIO X IDE: EXPERIMENTS. 



249 



apparatus so that both oxygen and sulphur dioxide can be passed 
together through a tuhp of hard ,-?lass as represented in Fig. 46. 




Fig. 46. 



First pass the two dried gases together through the empty tube 
and heat a part of the tube by means of a burner. Is there any 
evidence of combination ? Now stop the currents of the gases, 
let the tube cool down, and introduce a small layer of the plat- 
inized asbestos. Pass the dried gases over the heated asbestos. 
What takes place ? * 



CHAPTER XV. 
NITROGEN.— THE AIR.— ARGON, ETC. 

Nitrogen, N (At. Wt. 14.04). 

General. — Nitrogen bears to a group of elements rela- 
tions very similar to those which oxygen bears to the sul- 
phur group, and fluorine to the chlorine group. There 
are easily recognized resemblances between it and the 
members of the group, and yet there are some marked 
differences. As has been stated, and as is seen from its 
position in the periodic system, nitrogen is trivalent 
towards hydrogen, as shown in the compound NH 3 , while 
it is both trivalent and quinquivalent towards oxygen, as 
appears to be shown in N 2 3 and s N 2 5 . The principal 
hydrogen compound, ammonia, is entirely different in 
character from those of chlorine and sulphur, for, while 
these are acid, ammonia has in a marked way the character 
of a base, acting, however, in a peculiar way upon acids 
to form salts. The two oxides above referred to are acidic, 
forming the acids HN0 2 and HN0 3 , which are known as 
nitrous and nitric acids respectively. 

Occurrence of Nitrogen. — It was discovered by Lavoisier 
and Scheele towards the end of the last century that the 
air consists of two gases, one of which is oxygen, and they 
showed that when the oxygen is removed the gas which is 
left has not the power to support combustion nor to sup- 
port respiration. This gas was first called azote (from a, 
privative, and ^gotikos, life), and this name is still 
retained in France, the symbol in use in that country being 
Az, whereas in all others the symbol is N. This is the 
only case in which there is a difference of usage in respect 

250 



NITROGEN. 251 

to the symbols of the chemical elements in different 
countries. The name rritrogene was given to it later, from 
the fact that it is a constituent of nitre or saltpetre, KN0 3 
(nitrum, saltpetre, and ytveir, to produce), and this is 
the origin of the English name nitrogen. Not only is 
nitrogen found free in the air, but it is found in combina- 
tion in a large number of substances in nature. It is found 
in the nitrates, or salts of nitric acid, particularly as the 
potassium salt KN0 3 , and the sodium salt NaN0 3 , which 
occurs in enormous quantities in Chili, and is therefore 
known as Chili saltpetre. It is also found in the form of 
ammonia, which is a compound of nitrogen and hydrogen 
of the formula NH 3 . Ammonia occurs in small quantity 
in the air, and is formed under a variety of conditions, 
which will be taken up later. Nitrogen occurs, further, 
in combination in many animal substances. 

Preparation. — The most convenient way to prepare 
nitrogen is by burning in a closed vessel something which 
does not give a gaseous product of combustion; or by pass- 
ing air over something which has the power to unite with 
oxygen. The best substance to use for the first purpose is 
phosphorus, which burns readily and yields a solid product, 
soluble in water. It is only necessary, therefore, to place 
a piece of phosphorus in a floating vessel on the surface of 
water, set fire to it, and immediately place over it a closed 
bell-jar. As soon as the oxygen is used up the combustion 
stops, and the vessel then contains the residual nitrogen, 
and the walls are covered with a thin layer of phosphorus 
pentoxide, P 2 5 . This is soon converted by the water into 
phosphoric acid, which dissolves. Another convenient 
method for preparing nitrogen consists in passing air over 
copper heated in a tube. The copper takes up the oxygen 
readily, and the nitrogen passes on. Another good method 
consists in exposing to the air copper turnings partly 
covered with a solution of ammonia in a vessel so arranged 
as to allow free access of air while the escape of the gas in 
the vessel is prevented. This mixture absorbs oxygen 
slowly at the ordinary temperatures. Nitrogen can also 



25 2 COLLEGE CHEMISTRY. 

be made from other substances than the air. Thus, when 
chlorine is passed into a water solution of ammonia this 
reaction takes place: 

NH S + 3C1 = N + aHGLj 

but the hydrochloric acid combines at once with ammonia 
to form ammonium chloride, NH 4 C1: 

NH S + HC1 = NH 4 G1; 

so that the only gaseous product is nitrogen. This experi- 
ment is more or less dangerous, for if all the ammonia 
should be used up, and the passage of chlorine continued, 
a compound of nitrogen and chlorine which is extremely 
explosive is formed. Finally, nitrogen can be made by 
heating ammonium nitrite, NH 4 N0 23 either dry or in 
solution. The hydrogen and oxygen of the compound 
unite to form water and the nitrogen is set free: 

NH 4 NO a = 2H 2 + Ng. 

The nitrogen prepared from the air is never pure, as there 
are always present in the air other substances besides 
nitrogen and oxygen: and while some of these can be 
removed without serious difficulty, others cannot be. 

Properties. — Nitrogen is a colorless, tasteless, inodorous 
gas. At — 194° under the ordinary atmospheric pressure 
it is converted into a liquid: and this liquid solidities at 
— 214°. The critical temperature is — 146°, and the 
critical pressure is 35 atmospheres. A litre of nitrogen 
under standard conditions weighs 1.2505 grams. Its 
specific gravity (air = 1) is 0.967. It does not support 
combustion, nor does it burn. This latter fact is obvious, 
for, if nitrogen had the power to combine with oxygen 
when the temperature of the mixture is elevated, it is 
plain that this process of combustion would long ago have 
taken place, leaving one or the other of the two gases ami 
the product of combustion as the constituents of the air. 
Nitrogen not only does not combine with oxygen readily. 



NITROGEN. 253 

but it does not combine readily with any other element 
except at a very high temperature. Just as it does not 
support combustion, so also it does not support respiration. 
Animals would die in it, not on account of any active 
poisonous properties possessed by it, but for lack of oxygen. 
In the air it serves the useful purpose of diluting the 
oxygen. If the air consisted only of oxygen, all processes 
of combustion would certainly be much more active than 
they now are. What effect the continued breathing of 
oxygen would have upon animals it is impossible to say." 

The Air. — The atmosphere of the earth, commonly 
called the air, consists essentially of the two elements 
nitrogen and oxygen in the proportion of 79 volumes of 
nitrogen to 21 volumes of oxygen, or, by weight, of 77 
per cent of nitrogen and 23 per cent of oxygen. The 
weight of a litre of air under standard conditions is 1.293 
grams. Wherever air has been collected and analyzed it 
has been found to have practically the same composition. 
Nevertheless very accurate analyses have shown that the 
composition of the air is subject to slight variations. To 
decide whether the air is a chemical compound or a 
mechanical mixture requires a careful examination of a 
number of facts. The evidence may be summed up as 
follows : 

(1) If nitrogen and oxygen are mixed together the mix- 
ture conducts itself in exactly the same way as air. The 
mixing is not attended by any phenomena indicating 
chemical action. Generally the chemical combination of 
two elements is accompanied by an evolution of heat, and 
whenever a chemical act takes place there is some change 
in the temperature of the substances. When nitrogen and 
oxygen are brought together there is no change in the 
temperature of the gases, and no change of volume. 

(2) The law of definite proportions is founded upon a 
very large number of observations, and in all cases in 
which we have independent evidence that chemical action 
takes place it is found that the substances combine in 
exactly the same proportions to form the same product. 



2 54 COLLEGE CHEMISTRY. 

Variation in the composition of a chemical compound is 
not known. The composition of the air varies slightly, 
according to circumstances, and this fact may be regarded 
as evidence that the air is not a chemical compound. 

(3) Air dissolves somewhat in water. If air which is in 
solution in water is pumped out and analyzed, it is found 
to have a different composition from that of ordinary air. 
Instead of containing nearly 4 volumes of nitrogen to 1 of 
oxygen, it contains only 1.87 volumes of nitrogen to 1 of 
oxygen. The proportion of oxygen is much larger in the 
air which has been dissolved in water than it is in ordinary 
air. This is due to the fact that oxygen is more soluble in 
water than nitrogen. Therefore, when air is shaken with 
water, relatively more oxygen than nitrogen is dissolved. 
If the gases were in chemical combination we should 
expect the compound to dissolve as such and without 
change of composition. 

The above evidence shows that nitrogen and oxygen are 
not combined chemically in the air, but that they are 
simply mixed together. 

Analysis of Air. — The earliest examinations of Priestley, 
Lavoisier, and Scheele were made by burning substances 
in air contained in closed vessels. They concluded that 
the air is made up of \ oxygen and -£ nitrogen by bulk. 
In order to determine the composition of the air to-day, 
we should proceed as follows: A qualitative examination 
would easily show the presence of nitrogen and oxygen. 
If a solution of calcium ' hydroxide, Ca(OH) 2 , which is 
known as lime-water, or a solution of barium hydroxide, 
Ba(OH) 2 , is exposed to the air it becomes turbid, and a 
precipitate is formed. Neither nitrogen nor oxygen nor 
an artificially prepared mixture of the two gases can pro- 
duce this change. It has been shown that the change is 
due to the presence in the air of a small quantity of the 
gaseous compound, carbon dioxide, C0 2 . If calcium 
chloride or phosphorus pentoxide is exposed to the air it 
soon becomes moist and after a time turns liquid. This 
effect has been found to be due to the presence of water- 



ANALYSIS OF AIR. 255 

vapor in the air. Nitrogen obtained from the air by 
passing the latter over copper, which abstracts the oxygen, 
has been shown to contain a small amount of an extremely 
inert gas, argon (which see). By other methods which 
need not be considered here it can be shown that there are 
many other substances in the air besides those mentioned. 
Among them are ammonia, hydrogen dioxide, and organic 
matters of various kinds, including a large variety of 
germs the presence of which can be detected by the 
changes which exposure to the air produces in certain 
liquids, as milk and fruit- juices. 

Having thus learned what the chief constituents of the 
air are, the next thing is to determine in what quantities 
they are present, or to make a quantitative analysis of the 
air. For this purpose advantage may be taken of the fact 
that phosphorus when exposed to the air at ordinary tem- 
peratures combines slowly with the oxygen, leaving the 
nitrogen, argon, and minute quantities of other gases. 
If, therefore, a piece of ordinary phosphorus is inserted 
into a measured volume of air contained in a graduated 
glass tube over water or mercury, a diminution in volume 
will take place slowly. If, in the course of a few hours, 
the volume is again measured, the difference will give the 
volume of oxygen absorbed, while the gas remaining is 
nitrogen. Of course, in this case as in all others in which 
gas volumes are measured, corrections for temperature, 
pressure, and tension of aqueous vapor must be made. 

Another method by which the ratio between the nitrogen 
and oxygen in air can be determined is that which was 
first employed by Dumas and Boussingault. It consists in 
passing air over heated copper, collecting and measuring 
the nitrogen, and weighing the copper oxide. The appa- 
ratus is arranged as shown in Fig. 47. 

The copper is contained in the glass tube ab on the 
combustion furnace. At the ends of this tube are the 
stop- cocks rr. V is a glass globe provided with a stop- 
cock u. Before the experiment the air is exhausted from 
the globe and the tube ab, and the tube then carefully 



256 



COLLEGE CHEMISTRY. 



weighed. The tubes B and C and the apparatus A con- 
tain substances which have the power to absorb the carbon 
dioxide of the air: The tube ab is now heated and air 
admitted after passing through C, B, and A. The copper 




Fig. 47. 



takes up the oxygen, and the nitrogen enters the globe V. 
After the globe is full it is weighed, then exhausted and 
weighed again, and the difference gives the weight of the 
nitrogen. The tube is also exhausted and weighed, and 
the difference between this weight and that of the ex- 
hausted tube before the experiment gives the weight of the 
oxygen. 

The most refined method for the analysis of the air is 
the eudiometric method of Bunsen. This consists in 
adding some pure hydrogen to a measured volume of air 
contained in a eudiometer (see p. 74) over mercury, and 
then exploding the mixture by means of an electric spark. 
If the conditions are right all the oxygen present will 
combine with hydrogen, and in consequence of this there 
will be a corresponding contraction in the volume of the 
gases. The amount of contraction will be equal to the 
volume of hydrogen and that of oxygen which have com- 
bined to form water. But we know from previous experi- 
ments on these two gases that they combine in the ratio 



ANALYSIS OF AIR. 257 

of one volume of oxygen to two of hydrogen. Conse- 
quently the volume of oxygen which was present is equal 
to one-third of the total contraction. Of course it is 
necessary that there should be enough hydrogen present 
to combine with all the oxygen. This method is capable 
of great exactness. The most accurate analyses made by 
this method by Bunsen and others have shown that in 100 
volumes of air there are 20.9 to 21 volumes of oxygen. 

The estimation of the quantity of water-vapor present 
in the air is an important problem. The quantity present 
depends upon a variety of causes, the temperature and the 
direction of the wind being the chief ones. A good 
chemical method for estimating the water consists in draw- 
ing a known volume of air over calcium chloride in a 
weighed tube. This substance has the power to take up 
water, as we have repeatedly seen. If the tube is weighed 
after a certain volume of air has been drawn through it, 
the increase in weight will show the weight of water con- 
tained in that volume of air. 

The quantity of water-vapor present in the air varies 
between comparatively wide limits. At any given tem- 
perature the air cannot hold more than a certain quantity. 
When it contains this quantity it is said to be saturated. 
If cooled down below this temperature the vapor partly 
condenses, and appears as water. When a vessel contain- 
ing ice is placed in the air, that which immediately sur- 
rounds the vessel is cooled down below the point at which 
the quantity of water-vapor present would saturate the air, 
and water condenses on the outside of the vessel. Every 
one has noticed that on a warm cloudy day more water 
condenses on such a vessel than on a clear cool day. The 
water-vapor present in the air has an important effect on 
man. The inhabitants of countries with moist climates 
apparently have characteristics which are not generally 
met with in those who inhabit countries with dry climates. 
The difference between the effect of moist and that of dry 
air on an individual is well known. 

When air which is charged with water-vapor comes in 



258 COLLEGE CHEMISTRY. 

contact with cooler air, the vapor condenses and falls as 
rain. 

The method employed for the purpose of estimating the 
quantity of carbon dioxide in the air consists in drawing 
a known volume of air over something that has the power 
to absorb the carbon dioxide, and then determining the 
increase in weight of the absorbing substance. Potassium 
or sodium hydroxide is well adapted to this. An apparatus 
has been constructed in which barium hydroxide, Ba(OH) 2 , 
is used as the absorbent. When carbon dioxide is passed 
through a solution of this substance insoluble barium car- 
bonate, Ba00 3 , is thrown down according to the equation 

Ba(OH) 2 + C0 2 = BaC0 3 + H 2 0. 

This may be filtered off and weighed, and the quantity of 
carbon dioxide estimated from the results; or, if a known 
quantity of the hydroxide is taken, the quantity left 
unacted upon after the experiment can be determined by 
neutralizing with an acid, the neutralizing power of which 
has previously been determined with care. The quantity 
of carbon dioxide present in the air is relatively very small, 
being about 3 parts in 10,000. It varies slightly according 
to the locality and season, being greater in cities and in 
summer than in the country and in winter: and greater in 
warm countries than in cold. It is as essential to the life 
of plants as oxygen is to the life of animals. 

It is not an easy matter to determine the quantities of 
the other constituents of the air, as the ammonia, organic 
substances, etc., though there is no difficulty in determin- 
ing that they are present in very small quantities. 

Air and Life. — The relations of the air to the most im- 
portant chemical changes which are taking place upon the 
earth form one of the most interesting subjects of study. 
We have had a slight glimpse of the action of oxygen, and 
of that of carbon dioxide; both are essential to the life of 
plants and animals. So, too, the water- vapor acts chemic- 
ally upon plants, and probably to some extent in the 
respiration of animals. As regards the nitrogen, this 



PURE AND IMPURE AIR. 259 

element is frequently referred to as inert, and as serving 
the purpose of diluting the oxygen. Inert it undoubtedly 
is, and there is also no doubt that it dilutes the oxygen, 
but these statements give a very inadequate conception of 
the important part played by it in the processes of nature. 
Nitrogen in some form of combination is an essential con- 
stituent of plants and animals. The animals get their 
nitrogenous compounds from the plants, and the plants 
get theirs partly at least from the soil. By the growth of 
plants, therefore, nitrogenous compounds are constantly 
being withdrawn from the soil. When plants and animals 
undergo decomposition in the soil, the nitrogen contained 
in them is gradually converted into salts of nitric acid or 
nitrates, and if the decomposition takes place in the air 
the nitrogen is converted principally into ammonia. Both 
in the form of nitrates and of ammonia the nitrogen can 
be utilized by plants, so that if the plants and animals 
which have received their nourishment from a certain tract 
of land should be allowed to decay upon this land after 
death, and the products thus formed should be uniformly 
distributed in the soil, the latter would not become 
exhausted. But the products of the soil are removed, 
and, therefore, the nitrogen required for the growth of 
other plants is removed, and the soil becomes unproduc- 
tive. In order that it may be rendered fertile again, the 
lost nitrogen must be supplied. It appears from recent 
very elaborate experiments that the plants have the power 
to take up from the air a part of the nitrogen which they 
need. It has been shown that in this absorption of 
nitrogen from the air certain minute organisms that exist 
in the soil play a part. 

Pure and Impure Air. — Pure air may be denned as air 
that consists of nitrogen, oxygen, and carbon dioxide in 
the proportions stated above, together with some water- 
vapor, ammonia, hydrogen dioxide, and argon, and noth- 
ing of an injurious nature. It is evident from what has 
been said that there is constant danger of contamination 
from natural causes. The most common cause of con- 



260 COLLEGE CHEMISTRY. 

tamination is the breathing of human beings in rooms 
which are inadequately supplied with air. The breathing 
process involves the using up of oxygen and the giving on 
of carbon dioxide and small quantities of organic matter 
which is undergoing decomposition. If the quantity of 
oxygen is reduced below a certain limit the air becomes 
unfit for breathing purposes, and evil effects follow. An 
ordinary inhalation will not then be sufficient to supply the 
blood with the oxygen necessary to purify it, and the 
system will begin to suffer. Headache, drowsiness, and a 
general sense of discomfort follow. The ill effects of 
breathing the air of a badly-ventilated room occupied by a 
number of human beings appear to be due for the most 
part to the presence of the small quantities of decomposing 
organic matters which are given off from the lungs with 
the carbon dioxide and other gases. These act as poisons. 
They have been thrown off from the lungs because they 
are unfit for use, and when they are taken back again the 
normal processes of the body are interfered with. The 
subject of ventilation has been so thoroughly discussed of 
late years that great improvement has been made in the 
arrangements for supplying pure air to dwelling apart- 
ments and audience halls, but there is still room for 
improvement. Fortunately, in most buildings there is 
one source of supply of pure air which is independent of 
architects' plans. This is the diffusion of gases through 
the porous materials of which the buildings are con- 
structed. This diffusion was referred to under the head 
of Hydrogen (see p. 58). There is also a good deal of 
ventilation through the cracks and other apertures which 
are always to be found in our buildings. 

Under some conditions not thoroughly understood the 
air becomes contaminated by the decomposition of animal 
and vegetable matter. Air thus contaminated may cause 
specific diseases, and some of these are spoken of as being 
caused by malaria, a word which signifies simply bad air. 
It has been shown that there are in the air microscopic 
germs which have the power to develop in the body, and 



ARGON. 261 

then to cause the symptoms which are referred to malaria. 
Disease germs and germs of other kinds are present in the 
air in great variety, and they play an important part in 
connection with the life and health of mankind. 

Liquid Air. — Attention has been called to the fact that 
both oxygen and nitrogen can be liquefied. As air is a 
mixture of these two gases it follows that it can be liquefied. 
In brief the method employed consists in subjecting the air 
to a pressure of two thousand to twenty-five hundred 
pounds to the square inch and cooling this compressed air 
down to the ordinary temperature by water. It is then 
passed- through a tube that ends in a needle-valve. It is 
allowed .to escape through this valve into another tube that 
surrounds the one from which it has just escaped. In 
expanding it is cooled down below the temperature at 
which the compressed air in the inner tube becomes liquid. 
This liquid escapes with the gaseous air and is easily col- 
lected. 

Liquid air is a turbid, colorless liquid. The turbidity 
is due to the presence of solid water and solid carbon 
dioxide. By passing the liquid through a paper filter the 
solids are removed, and a transparent liquid is thus 
obtained. This consists mostly of nitrogen and oxygen in 
the proportion of about four-fifths of the former to one- 
fifth of the latter. This mixture boils at — 191°. As the 
nitrogen boils at a lower temperature (— 194°) than 
oxygen (— 181°), more nitrogen than oxygen is converted 
into gas in a given time, and after a time the liquid that 
is left is much richer in oxygen than ordinary air. 

Argon. — Lord Kayleigh has shown that a litre of nitrogen 
prepared from the air by abstracting the oxygen and the 
small quantities of other substances known to be present 
weighs 1.2572 grams, while a litre of nitrogen made from 
some chemical compound, such as ammonia, weighs 1.2505 
grams. Chemically-prepared nitrogen is lighter than that 
obtained from the air. This observation led Lord Rayleigh 
and W. Ramsay to a more thorough chemical examination 
of the air, the result of which was the discovery of a con- 



262 COLLEGE CHEMISTRY. 

stituent previously unknown. This is the gas argon. It 
can be obtained by either of two methods: (1) By passing- 
air over heated copper until all the oxygen is abstracted, 
and then over heated magnesium which unites with the 
nitrogen but leaves the argon; (2) by mixing air with 
oxygen and passing electric sparks through the mixture. 
The oxygen and nitrogen combine and the product can 
easily be removed. After all the nitrogen has thus been 
removed argon remains behind. Argon is an element with 
the atomic weight about 40. It is present in the air to 
the extent of about 1 per cent of the nitrogen. It cannot 
be made to combine with any other element. 

Other Gases in the Air. — By allowing the greater part 
of liquid air to evaporate and examining the residue 
Eamsay has discovered another gaseous substance in the 
air which he calls krypton. This has since been obtained 
in larger, and probably purer, condition by Ladenburg 
and Kruegel. It has the specific gravity 58.67, or about 
59 (H = 1). 

By liquefying argon by means of liquid air, and collect- 
ing the gas given off from it at successive stages, Eamsay 
has shown, further, that two other gases, which he calls 
neon and xenon, are present in the air in very small 
quantities. 

. EXPERIMENTS. 

Preparation of Nitrogen. 

Experiment 115. — Place a good-sized stoppered bell-jar over 
water in a pneumatic trough. In the middle of a flat cork about 
three inches in diameter fasten a small porcelain crucible, and 
place this on the water in the trough. Put in it a piece of phos 
phorus about twice the size of a pea, and set fire to it. Quickly 
place the bell-jar over it. At first some air will be driven out of 
the jar. The burning will continue for a short time, and then 
gradually grow less and less active, finally stopping. On cooling, 
it will be found that the volume of gas is less than four-fifths the 
original volume, for the reason that some of the air was driven 
out of the vessel at the beginning of the experiment. Before re- 



EXPERIMENTS: ANALYSIS OF AIR. 



263 



moving the stopper of the bell-jar see that the level of the liquid 
outside is the same as that inside. Try the effect of introducing 
successively several burning 
bodies into the nitrogen, — as, 
for example, a candle, a piece 
of sulphur, phosphorus, etc. 

Experiment 116.— Pass air 
slowly over copper contained 
in a tube heated to redness 
and collect the gas which 
passes through. Does it act 
like nitrogen ? 

Experiment 117.— In a good- 
sized WoulfFs bottle provided 
with a safety-funnel and de- 
livery-tube as shown in Fig. 
48, put some copper-turnings 
and pour upon them concen- 
trated ammonia, but not 
enough to cover them. Close 
the delivery-tube by means of 
a pinch-cock ; and let the ves- 
sel stand, What evidence of 
action is there ? After a time, 
force some of the gas out of 
the bottle by pouring water 

through the funnel, and opening the delivery-tube, 
gas act like nitrogen ? 




Fig. 43. 



Does the 



Analysis of Air. 

Experiment 118. — Arrange an apparatus as in Fig. 11. In- 
stead of a plain tube, use one graduated in cubic centimetres. 
Enclose 60 to 80 cc. air in the tube over water. Arrange the 
tube so that the level of the water inside and outside is the same. 
Note the temperature of the air and the height of the barometer. 
Reduce the observed volume to standard conditions. Now intro- 
duce a piece of phosphorus, as in Experiment 26, and allow it to 
stand for twenty-four hours. Draw out the phosphorus. Again 
arrange the tube so that the level of the water inside is the same 
as that outside. Make the necessary corrections for temperature, 
pressure, and the pressure of aqueous vapor. It will be found 
that the volume has diminished considerably, but that about four 
fifths of the gas originally put in the tube is still there. If the 



264 COLLEGE CHEMISTRY. 

work is done properly, the volume of the gas left in the tube will 
be to the total volume used as 79 to 100. In other words, of 
every 100 cc. air used 21 cc. are absorbed by phosphorus, and 
79 cc. are not. The gas absorbed is oxygen, identical with the 
oxygen made from the oxide of mercury, manganese dioxide, 
and potassium chlorate. The gas left over has no chemical prop- 
erties in common with oxygen. Carefully take the tube out of the 
vessel of water, closing its mouth with the thumb or some suit- 
able object to prevent the contents from escaping. Turn it with 
the mouth upward, and introduce into it a burning stick. Does 
it support combustion ? Is it oxygen ? 

Experiment 119. — Expose a few pieces of calcium chloride on 
a watch-glass to the air. It gradually becomes liquid by absorb- 
ing water from the air. 

Experiment 120. — Expose some clear lime-water to the air. It 
soon becomes covered with a white crust. A similar change 
takes place if baryta-water is exposed in the same way. Lime- 
water is made by putting a few pieces of quicklime in a bottle 
and pouring water upon it. The mixture is well shaken up and 
allowed to stand. The undissolved substance settles to the bot- 
tom, and with care a clear liquid can be poured off the top. 
This is lime-water, which is a solution of calcium hydroxide, 
Ca(OH) 3 , in water. Baryta-water is a solution of a similar com- 
pound of the element barium. When these solutions are exposed 
to nitrogen or oxygen, or to an artificially prepared mixture of 
the two gases, no change takes place. Further, if air is first 
passed through a solution of caustic soda it no longer has the 
power to cause the formation of a crust on lime-water or baryta- 
water. 

Experiment 121. — Arrange an apparatus as shown in Fig. 49. 
The wash- cylinders J. and B are half filled with a solution of ordi- 
nary caustic soda. The bottle C is filled with water. The tube 
J), which should be filled with water and provided with a pinch- 
cock, acts as a siphon. Open the pinch-cock and let the water 
flow slowly out of the bottle. As it flows out air will be drawn 
in through the caustic soda in the wash-cylinders. When the 
bottle is a quarter filled with air pour some water in again until 
it is full. Then draw all the w r ater off. Now remove the stopper 
from the bottle, pour in 20 to 30 cc. lime-water and cork the 
bottle. The crust formed on the lime-water will now be hardly, 
if at all, perceptible. There is, therefore, something present in 
the air under ordinary circumstances which has the power to 
form a crust on lime-water or baryta-water, and which can be 



EXPERIMENTS : ANALYSIS OF AIR. 



265 



removed by passing the air through caustic soda. Thorough 
examination has shown that this is the compound which chemists 









' — "^~^U~^ ~ 


"1 






^•/fI s^ 




















(7 




kfW 






-^l 


B 




J 


















y 


V 

> 


i— 


u 







t 



Fig. 49. 



call carbon dioxide, and which is commonly known as carbonic 
acid gas. It is the substance which was obtained by burning 
charcoal in oxygen. 




Fig. 50. 



Experiment 122. — Into the bottle containing the air from 
which the carbon dioxide has been removed hold a burning stick 



266 COLLEGE CHEMISTRY. 

or taper for a moment. Notice whether a crust is now formed 
on the lime-water. Wood and the material from which the taper 
is made contain carbon. Explain the formation of the crust on 
the lime-water after the stick of wood or taper has burned for a 
short time in the vessel. 

Experiment 123. — Arrange an apparatus as shown in Fig. 50. 
The bottle A contains air ; B contains concentrated sulphuric 
acid ; G contains granulated calcium chloride ; D is carefully dried 
and contains a few pieces of granulated calcium chloride and 
air. Pour water through the funnel-tube into A, when the air 
will be forced through B and C and into D. But in passing 
through B and C the moisture contained in it will be removed, 
and the air which enters D will be dry. After A has once been 
filled with water, empty it and fill it again, letting the dried air 
pass into D. This operation may be repeated any number of 
times. The calcium chloride in 1) will not grow moist; 



CHAPTER XVI. 

COMPOUNDS OF NITROGEN WITH HYDROGEN.— 
WITH HYDROGEN AND OXYGEN.— WITH OXYGEN, 
ETC. 

General Conditions which give Rise to the Formation 
of the Simpler Compounds of Nitrogen. — We have seen 
that at ordinary temperatures nitrogen is an inactive 
element, showing little tendency to combine with other 
elements. It is nevertheless an easy matter to get com- 
pounds of nitrogen with many other elements, and among 
these compounds, some of those which it forms with 
hydrogen and oxygen are of high importance. 

When a compound that contains carbon, hydrogen, and 
nitrogen, and is not volatile, is heated in a closed vessel, so 
that the air does not have access to it, the nitrogen passes 
out of the compound, not as nitrogen, but partly in com- 
bination Avith hydrogen, in the form of the compound 
ammonia. Nearly all animal substances contain carbon, 
hydrogen, oxygen, and nitrogen in many forms of com- 
bination, some of which are quite complicated. Many 
of these give off ammonia when heated. Similarly, com- 
pounds containing carbon, oxygen, and hydrogen, even 
though they may be thoroughly dry, when heated give off 
oxygen in combination with hydrogen in the form of 
water. Both these kinds of decomposition, that which 
gives ammonia and that which gives water, are to be 
ascribed to the fact that the compounds of carbon which 
are heated are unstable at higher temperatures, and when 
they are broken down the elements contained in them 
arrange themselves in combination in stable forms, such 
as the comparatively simple compounds, water and am- 

267 



268 COLLEGE CHEMISTRY. 

monia. An illustration of this kind of action was referred 
to in speaking of the preparation of nitrogen by heating 
ammonium nitrite. This compound breaks down very 
easily under the influence of heat, the hydrogen and 
oxygen combining to form the stable compound water, 
while the nitrogen remains uncombined. Some animal 
substances, as, for example, urine, give off ammonia when 
they undergo spontaneous decomposition in the air. This 
decomposition is generally, if not always, due to the action 
of minute organisms, the germs of which are in the air, 
which develop when they come in contact with certain 
substances. The coal which is used for making illuminat- 
ing gas contains some hydrogen and- nitrogen in chemical 
combination, and when the coal is heated ammonia is given 
off with the other products. 

When animal substances undergo decomposition in the 
presence of basic compounds where the temperature is 
comparatively high, the nitrogen combines with oxygen 
and with the metal of the base. Either a salt of nitric 
acid, HN0 3 , or of nitrous acid, HN0 2 , is formed. In 
some countries where the conditions are favorable to the 
process, immense quantities of nitrates are found, chiefly 
potassium nitrate, KN0 3 , and sodium nitrate, NaN0 3 . 
Nitrates are, however, found everywhere in the soil. The 
change of animal and vegetable nitrogenous substances to 
the form of nitrates is caused by the action of minute living 
organisms, which are found everywhere, and serve an im- 
portant purpose in converting the waste animal and vege- 
table matter into simple compounds that can be utilized by 
plants. How they effect the change is not known. From 
the salts of nitric acid which are found in nature, nitric 
acid itself can easily be made. 

Nearly all the compounds of nitrogen with which we 
have to deal are made either from ammonia or from nitric 
acid. 

Relations between the Principal Compounds of Nitro- 
gen. — In studying the compounds of sulphur, we saw that 
whenever a compound of sulphur is oxidized with a strong 



COMPOUNDS OF NITROGEN. 269 

oxidizing agent the final product of the action is sulphuric 
acid; and that, on the other hand, under the influence of 
strong reducing agents these compounds yield hydrogen 
sulphide as the final product. So, also, the limit of oxidiz- 
ing action in the case of chlorine is perchloric acid, and 
of reducing action, hydrochloric acid. Most of the other 
compounds of sulphur with hydrogen and oxygen are 
products intermediate between the two limits, sulphuric 
acid and hydrogen sulphide, as the other compounds < f 
chlorine with hydrogen and oxygen are intermedial e 
between hydrochloric acid and perchloric acid. The limit 
of reduction of nitrogen compounds is ammonia, NH 3 , 
and of oxidation, nitric acid, HN0 3 . As has been noticed, 
the valence of nitrogen towards oxygen is greater than it 
is towards hydrogen, but the difference is not as marked 
as in the case of the members of the chlorine group, and 
in that of the sulphur group. Its hydrogen valence is 3, 
its maximum oxygen valence is 5. 

A tabular list of the oxides of nitrogen and the acids 
formed by it is here given: 

Oxides. Acids. 

Nitrous oxide, N 2 Hyponitrous acid, N 2 (OH) 2 

Nitric oxide, NO Nitrous acid, NO (OH) 

Nitrogen trioxide, | ^ ~ Nitric acid, N0 2 (OH) 
(Nitrous anhydride), f 2 3 
Nitrogen peroxide, N0 2 (N 2 4 ) 
Nitrogen pentoxide, ) ■«■ n 
(Nitric anhydride),- j" ^^ 

Besides the above there are the basic compounds, 
hydroxylamine, NH 2 (OH), ammonia, NH 3 , and hydrazine, 
N 2 H 4 ; and with water ammonia probably forms the hy- 
droxide NH 4 (OH), which is a base. In addition to these 
there is a compound of nitrogen with hydrogen of the 
formula N 3 H. This is an acid. (See Triazoic Acicj.) 
There is another acid of the same composition as hypo- 
nitrous acid but differing from it in properties, the rela- 
tion of which to hyponitrous acid is not yet quite clear. 



270 COLLEGE CHEMISTRY. 

With the members of the chlorine group nitrogen forms 
a few compounds which are characterized by marked 
instability. So unstable are they that they explode violently 
when simply touched. The chloride of nitrogen explodes 
with great violence when the direct rays of the sun are 
allowed to shine upon it. With sulphur, nitrogen forms 
two compounds. With sulphur, hydrogen, and oxygen, 
however, nitrogen forms a number of compounds, one of 
which has already been referred to in connection with the 
manufacture of sulphuric acid. This is the so-called 
nitrosyl-sulphuric acid which is formed by the action of a 
mixture of the peroxide, N0 2 , and the monoxide, NO, on 
sulphurous acid, oxygen, and water. 

In studying the compounds of nitrogen, it will be best 
to begin with the end products, ammonia and nitric acid. 

Ammonia, NH 3 . — The conditions under which ammonia 
is formed have been mentioned. The chief source at 
present is the "ammoniacal liquor" of the gas-works. 
This is the water through which the gas has been passed 
for the purpose of removing the ammonia, and it contains 
ammonia, or ammonium hydrc \I le, NH 4 (OH), in solution. 
By adding hydrochloric acid to this solution the salt 
ammonium chloride, NH 4 C1, is formed. This is the well- 
known substance sal ammoniac. It appears that this name 
has its origin in the fact that common salt or sodium 
chloride. NaCl, was formerly called sal armeniacum, and 
that afterward, through a misunderstanding, 'immonium 
chloride came to be known by the same name which under- 
went change to the form sal ammoniacum, or sal ammoniac. 
When the ammoniacal liquor is treated with sulphuric acid, 
ammonium sulphate is formed. From one or the other of 
these salts it is a simple matter to obtain ammonia. For 
this purpose it is only necessary to treat the salt with some 
strongly basic compound, as, for example, potassium or 
sodium hydroxide, or calcium hydroxide. Thus, when a 
solution of potassium hydroxide is poured on ammonium 
chloride or sulphate the strong penetrating odor of am- 



AMMONIA. 271 

monia is at once noticed. The first reaction probably 
results in the formation of ammonium hydroxide, thus: 

NH 4 C1 + KOH = NH 4 OH + KC1; 
(NH 4 ) 2 S0 4 + 2KOII = 2NH 4 OH + K 2 S0 4 ; 
2NH 4 C1 + Ca(OH) 2 = 2NH 4 OH + CaCl 2 ; 
(NH 4 ) 2 S0 4 + Ca(OH) 2 = 2NH 4 OH + CaS0 4 . 

But the ammonium hydroxide breaks down very readily 
into water and ammonia, which escapes as a gas: 

NH 4 OH = NH 3 + H 2 0. 

In the laboratory ammonia is prepared by treating am- 
monium chloride with slaked lime or calcium hydroxide. 
The two are mixed in the proportion of two parts of slaked 
lime to one of ammonium chloride, placed in a flask and 
gently heated, when the ammonia is given off at once. 

It is frequently more convenient to heat a strong aqueous 
solution of ammonia, such as is found in every chemical 
laboratory. Such a solution when gently heated readily 
gives off ammonia. 

Properties. — Ammonia is a colorless, transparent gas 
with a very penetrating, characteristic odor. In concen- 
trated form it causes suffocation. Its specific gravity is 
0.59; that is to say, it is but little more than half as heavy 
as air. A litre of the gas under standard conditions weighs 
0.7635 gram. It can easily be compressed to the liquid 
form by pressure and cold. When the pressure is removed 
from the liquefied ammonia it passes back to the form of 
gas, and in so doing it absorbs a great deal of heat. These 
facts are taken advantage of for the artificial preparation 
of ice. This application will be clear from the following- 
explanation and Fig. 51. 

An aqueous solution of ammonia saturated at 0° is 
brought into the strong iron cylinder A, and then gently 
warmed, while the vessel B is cooled by cold water. The 
gas given off from A passes through the bent tubes into 
B, where it is condensed to a liquid. The cylinder A is 
now placed in a vessel of cold water, and the water which 



V 



272 



COLLEGE CHEMISTRY. 




Fig. 51. 



is to be frozen is placed in a cylinder D, and this into the 

\ollow space E in the vessel B. 
The liquid ammonia passes 
rapidly into the form of gas 
which is absorbed in the 
water in A, while at the same 
time so much heat is absorbed 
that the water in D is frozen. 
Ammonia does not burn 
in the air, but does burn in 
oxygen with a pale yellow- 
ish flame. It is absorbed by 
water in very large quantity. 
One volume of water at the 
ordinary temperature dissolves about 600 volumes of 
ammonia gas, and at 0° about 1000 volumes. The sub- 
stance with which we commonly have to deal under the 
name of ammonia is a solution of ammonia in water. It 
is called "spirits of hartshorn" in common language. 
The solution has the odor of the gas. It loses all its gas 
when heated to the boiling temperature. The solution 
shows a strong alkaline reaction, and has the power to 
neutralize acids and form salts. The conduct of the solu- 
tion is, in fact, strikingly like that of sodium and potassium 
hydroxides, and it is believed that in the solution there are 
present the ions of a compound of the formula NH 4 (OH), 
known as ammonium hydroxide, and formed by the direct 
action of ammonia upon water. If this is true, then the 
action of ammonia upon acids is to be explained as fol- 
lows: Ammonium hydroxide is analogous to potassium 
hydroxide, but differs from it in that it contains the group 
of atoms NH 4 in place of the atom K. In some way this 
group plays in the salts formed by ammonia the same part 
that the elementary atom potassium plays in the salts of 
potassium, and just as the latter are called potassium salts, 
so the former are called ammonium salts. According to 
this the ammonium salts are salts which contain the group 
NH^, known as the ammonium group, in place of the 



AMMONIA. 273 

hydrogen of the acids. They are formed by direct com- 
bination of ammonia with the acids, or by the action uf 
ammonium hydroxide upon acids. The analogy between 
the action of ammonium hydroxide and that of potassium 
hydroxide upon acids is clearly shown by the aid of the 
following equations : 

K(OH) +HC1 r=KCl +H 2 0; 

NH 4 (OH) +HC1 =NH 4 G1 + H a O: 
K(OH) ' + HNO s =KNO s + H 2 0; 
NH 4 (OH) + HN0 S = NH 4 N0, + H 2 0; 
2K(OH) -f H 2 S0 4 = K,S0 4 + 2H 2 0; 
2NH 4 (OH) + H 2 S0 4 = (NHJ,S0 4 + 2H 2 0. 

The formation of the ammonium salts by direct action 
of ammonia, NH 3 , upon acids is represented in the follow- 
ing equations: 

NH 3 + HOI = NH 4 C1; 
NH 3 4- HN0 3 = NH 4 N0 3 ; 

2NH 3 + H 2 S0 4 = (NH 4 ) 2 S0 4 . 

The strong tendency of ammonia to combine directly 
with acids is shown by bringing two uncovered vessels, one 
containing a solution of ammonia, and the other a solution 
of hydrochloric acid, near each other. A dense cloud will 
at once be formed, if the solutions are concentrated. This 
is due to the direct combination of the gases which escape 
from the solutions. 

While the assumption of the existence of the group 
ammonium, NH 4 , in the ammonium salts is of great 
service in dealing with these salts, and while this assump- 
tion appears to be entirely justified by the facts, no com- 
pound of this composition has as yet been isolated. The 
name ammonium is given to the hypothetical compound 
on account of the fact that it evidently plays the part of a 
metallic element, and it is customary to give such elements 
names ending in ium. While, further, it is generally 
believed that the ions of the compound ammonium 



2 74 COLLEGE CHEMISTRY. 

hydroxide, NH 4 (OH), are formed when ammonia dissolves 
in water, the compound itself has not been isolated, owing 
to its instability and tendency to break down into ammonia 
and water. On the other hand, some very interesting 
derivatives of this hydroxide have been isolated. There 
is one of these which is derived from the hydroxide by the 
replacement of the four hydrogen atoms of the ammonium 
by groups of carbon and hydrogen atoms. This compound 
is stable and can be isolated as a, hydroxide of the general 
formula NK 4 (OH), in which E 4 represents the groups of 
carbon and hydrogen atoms. In solution it acts almost 
exactly like potassium hydroxide. 

Composition of Ammonia. — By oxidation under the 
proper conditions it is possible to convert the hydrogen of 
ammonia into water and leave the nitrogen in the free 
state. As water and nitrogen are the only products formed, 
and the quantity of oxygen used up in the oxidation is 
equal to the quantity of oxygen found in the water formed, 
it follows that nitrogen and hydrogen are the only elements 
contained in ammonia. 

When electric sparks are passed for some time through 
a mixture of nitrogen and hydrogen, some ammonia is 
formed. Conversely, when electric sparks are passed for 
a time through ammonia, nitrogen and hydrogen are 
obtained. 

If, in the oxidation of a known weight of ammonia, the 
water formed and the nitrogen left uncombined are 
accurately determined, it will be found that in ammonia 
the elements are combined very nearly in the proportion 
of 14 parts by tueight of nitrogen to 3 parts ~by weight of 
hydrogen. Further, the molecular weight determined by 
the method of Avogadro is approximately 17. Therefore, 
the molecular formula of ammonia is NH 3 , the atomic 
weight of nitrogen being 14. 

The proportion by volume in which the two elements 
combine can be determined by the following method: A 
glass tube closed at one end and provided with a glass 
stojD-cock at the other is filled with pure chlorine gas. By 



COMPOSITION OF AMMONIA. 275 

means of a small funnel attached to the open end a little 
of a strong aqueous solution of ammonia is slowly intro- 
d.iced into the tube. Keaction takes place at once between 
the chlorine and the ammonia, according to the equation 

and the hydrochloric acid unites with ammonia to form 
ammonium chloride: 

3HC1 + 3NH 3 = 3NH 4 C1. 

The entire change is therefore represented by the equation 

4NH 3 + 301 = N + 3NH 4 C1. 

Hydrogen and chlorine combine in equal volumes, as w r e 
have already learned. Now, if we start with a measured 
volume of chlorine, and add ammonia to it until it is all 
used up, we know that the volume of hydrogen which has 
been extracted from ammonia is equal to the volume of 
chlorine with which we started. If we measure the volume 
of nitrogen left over, w r e know the volume of the nitrogen 
which was in combination with a volume of hydrogen equal 
to that of the chlorine originally taken. This experiment 
has been tried repeatedly, and it has been found that the 
ratio of the volume of nitrogen to that of the hydrogen 
with which it was combined is as 1 to 3. The tube being 
full of chlorine at the beginning of the experiment, it will 
be found to be one-third full of nitrogen at the end. 
Therefore, in ammonia 1 volume of nitrogen is combined 
luith 3 volumes of hydrogen. 

The experiment just referred to will perhaps be better 
understood by the aid of the accompanying diagram. The 
chlorine in the tube may be represented as made up of 
three equal parts or volumes. Each volume of chlorine 
combines with an equal volume of hydrogen, leaving the 
nitrogen uncombined. The volume of nitrogen left is only 



276 



COLLEGE CHEMISTRY. 



one-third of that of the chlorine, or for three volumes of 
chlorine there is one volume of nitrogen: 



1 vol. 
CI 




1vol. 
CI 




1vol. 
CI 



combine 
with 



1vol. 
H 




1vol. 
H 




1vol. 
H 



leaving 



vol. 

N 



Therefore in ammonia the gases nitrogen and hydrogen 
are combined in the proportion of 1 volume of nitrogen 
to 3 volumes of hydrogen: 



1 vol. 
H 



1vol. 
H 




1vol. 

N 



Volume relations in ammonia. 



lvol. 
H 



Another question in regard to the volume relations 
remains to be answered, and that is: When nitrogen and 
hydrogen unite in the proportions above stated, how many 
volumes of ammonia gas do the four volumes of the con- 
stituents form ? It is not possible to determine this by 
direct combination of the two gases, but ammonia can be 
decomposed into its constituents by continued passage of 
electric sparks through it. When this is done it is found 
that after the decomposition the gases occupy twice the 
volume which was occupied by the ammonia. It appears, 



AMMONIUM AMALGAM. 277 

therefore, that when hydrogen and nitrogen combine to 
form ammonia the volume is reduced to one-half, or, what 
is the same thing, when three volumes of hydrogen com- 
bine with one volume of nitrogen the four volumes form 
two volumes of ammonia gas. 

The above facts have already been commented upon in 
speaking of the combination of gases in general; and it 
has been shown that chlorine, oxygen, and nitrogen com- 
bine with hydrogen in entirely different ways : 

(1) 1 volume of chlorine combines with 1 volume of 
hydrogen to form 2 volumes of hydrochloric acid gas. 

(2) 1 volume of oxygen combines with 2 volumes of 
hydrogen to form 2 volumes of water-vapor. 

(3) 1 volume of nitrogen combines with 3 volumes of 
hydrogen to form 2 volumes of ammonia gas. 

What the cause of these differences is we do not know. 
In some way these facts are directly connected with the 
law of Avogadro that equal volumes of all gases contain 
the same number of molecules, and with the power of the 
atoms of chlorine, oxygen, and nitrogen to combine with 
one, two, and three atoms of hydrogen respectively. When 
one atom of chlorine unites with one atom of hydrogen the 
result is a molecule. So also when one atom of oxygen 
unites with two atoms of hydrogen, and when one atom 
of nitrogen unites with three atoms of hydrogen, the result 
in each case is a molecule, and, according to the law of 
Avogadro, a gaseous molecule, whether it consists of one 
atom or a hundred atoms, occupies the same space. 

Ammonium Amalgam. — A very curious substance which 
appears to consist of mercury and ammonium is formed 
when a solution of ammonium chloride is treated with a 
compound of sodium and mercury known as sodium 
amalgam. The action is thought to take place thus : 

2NH 4 C1 + Na 2 Hg = 2NaCl + (NH 4 ) 2 Hg. 

The product, ammonium amalgam, is unstable, break- 
ing down very soon into ammonia, hydrogen, and mercury. 



/ 



27S COLLEGE CHEMISTRY. 

It will be referred to again and somewhat more fully under 
the head of Mercury. 

Hydrazine, N 2 H 4 . — A compound closely related to am- 
monia and ammonium has recently been prepared by a 
complicated method that cannot be explained here. This 
is known as hydrazine. Its composition and molecular 
weight are represented by the formula N 2 H 4 . A large 
number of derivatives of hydrazine are known, and have 
been studied exhaustively. 

Hydrazine is a liquid that boils at 113.5° in a current of 
hydrogen. It acts upon acids much as ammonia does, 
forming the hydrazine salts. 

Hydroxylamine, NH 2 (OH). — This compound is pre- 
pared by reducing nitric acid : 

N0 2 (OH) _j_ 6H = NH 2 (OH) + 2H 2 0. 

It can also be prepared by reduction of nitric oxide : 

2NO + 6H = 2NH 2 (OH). 

It is a solid consisting of leaflets or hard needles. It 
melts at 33.05°, and boils at 58° under a pressure of 
22 mm. When the water solution is evaporated, both 
ammonia and hydroxylamine pass over with the water. 
Its salts are easily obtained by treating the solution with 
acids : 

NH 2 (OH) + HC1 = NH s (OH)Cl; 
NH 2 (OH) + HNO3 = NH 3 (OH)N0 3 . 

From the method of formation and the composition it 
appears that these salts are ammonium salts in which 
hydroxyl is substituted for one hydrogen of the ammonium. 
They should therefore be called hydroxyl- ammonium salts. 
One of the most characteristic properties of hydroxylamine 
is the ease with which it breaks down into ammonia, 
nitrogen, and water: 

3NH 2 OH = N 2 + NH 3 + 3H 2 0. 



V 
TRIAZOIC ACID— NITRIC ACID. 279 

If brought in contact with compounds capable of reduc- 
tion, it reduces them, the nitrogen in these cases generally 
combining with oxygen to form nitrous oxide, N 2 0. The 
reduction of cupnc oxide, CuO, takes place according to 
the equation 

2NH 2 (OH) + 4CuO = N 2 + 2Cu 2 + 3H 2 0. 

By nascent -hydrogen hydroxylamine is reduced to 
ammonia. 

Triazoic Acid, N 3 H. — This compound, which is also 
called lujdr azoic and hydro nitric acid, can be made by a 
number of reactions involving the use of complex organic 
substances. A simpler method consists in passing am- 
monia gas over heated metallic sodium, and then passing 
nitrous oxide, N 2 0, over the resulting product, which is 
sodium amide. The following reaction takes place : 

NH 2 Na + N 2 = N 3 Na + H 2 0. 

By dissolving in water the sodium salt thus formed, 
treating with dilute sulphuric acid, and distilling, a solu- 
tion of the acid is obtained. The compound is a colorless 
liquid that boils at 37°. It has a very penetrating odor, 
and produces bad effects upon one who inhales it. It is 
an acid resembling hydrochloric a.cid. It, as well as some 
of its salts, is extremely explosive. Its ammonium salt 
formed by direct union with ammonia is interesting as it 
has the composition KNH 4 or N 4 H 4 . 

Nitric Acid, HN0 3 . — This important chemical compound 
was first made, though not in pure condition, about the 
ninth century by distilling saltpetre, copper sulphate, and 
alum. The name nitric acid has its origin in the fact that 
the compound is formed from nitre. It has already been 
stated that the salts of this acid, particularly the potassium 
and sodium salts, occur very widely distributed in the earth, 
and that there is a great accumulation of the sodium salt 
in South America, whence the name Chili saltpetre. 
Wherever organic matter, particularly that of animal 



280 COLLEGE CHEMISTRY. 

origin, undergoes spontaneous decomposition in the pres- 
ence of basic substances, nitrates are formed, in conse- 
quence of the action of an organism known as the nitrifying 
ferment. This process of nitrification has already been 
referred to in a general way. It is one of great importance 
for the welfare of the human race, and indeed of most 
living beings, as by its aid the useless nitrogenous sub- 
stances of dead plants and animals are converted into the 
useful nitrates which in the soil aid the process of plant 
growth. 

Nitric acid can be formed by the action of electric 
sparks on nitrogen and oxygen in the presence of water. 
It is also formed by the action of oxidizing agents on 
ammonium compounds. 

Nitric acid is always prepared by treating potassium or 
sodium nitrate with concentrated sulphuric acid. When 
sodium nitrate is treated with sulphuric acid action takes 
place thus : 

NaN0 3 + H 2 S0 4 = NaHS0 4 + HN0 3 . 

The salt formed in this way is primary or acid sodium 
sulphate. If sufficient of the saltpetre is present and the 
temperature is raised, a second reaction takes place, result- 
ing in the formation of normal sodium sulphate : 

NaN0 3 + NaHS0 4 = Na 2 S0 4 + HN0 3 . 

But the temperature required for this reaction is so high 
that a considerable part of the nitric acid is decomposed. 
In the preparation of nitric acid, therefore, the first 
reaction is the one used, and for this purpose the sub- 
stances are brought together in retorts in the proportion 
of their molecular weights (about equal weights), and the 
retort gently heated. The nitric acid distils over slowly, 
and is condensed by cooling the receiver. 

On the large scale the acid is made by bringing Chili 
saltpetre and concentrated sulphuric acid together in cast- 
iron cvlinders or retorts. 



NITRIC ACID. 281 

Nitric acid is a colorless volatile liquid. It begins to 
boil at 86°, but at this temperature it undergoes partial 
decomposition into nitrogen peroxide, water, and oxygen : 

2HN0 3 = 2N0 2 + H 2 + 0. 

It undergoes the same change slowly when exposed to the 
direct rays of the sun. In consequence of this decomposi- 
tion the distillate collected in the manufacture of nitric 
acid, and, in general, whenever the acid is distilled, always 
contains a considerable percentage of water, and is colored 
more or less yellow by the nitrogen peroxide present. In 
order to abstract the water from ordinary nitric acid it is 
mixed with concentrated sulphuric acid and slowly dis- 
tilled; but even under these circumstances the product is 
colored in consequence of some decomposition, and it also 
contains some water. By conducting carbon dioxide gas 
through the gently warmed acid the nitrogen peroxide'can 
be removed, and in this way an acid containing about 99.5 
per cent of the compound HN0 3 has been obtained. 

Pure nitric acid is a very active substance chemically. 
It gives up its oxygen readily and is itself thus reduced to 
other compounds of nitrogen and oxygen, or of nitrogen, 
oxygen, and hydrogen, as has already been pointed out. 
When it acts upon the metals it forms nitrates, metal atoms 
being substituted for the hydrogen. According to the 
conditions, nitrogen peroxide, NO,, nitrous acid, HN0 2 , 
nitric oxide, NO, nitrous oxide N 2 0, nitrogen, hydroxyl- 
amine, Nil, (OH), and, finally, ammonia are formed by 
reduction of the acid. Of these reactions, that which gives 
nitric oxide, NO, is the one which commonly takes place 
on treating metals with nitric acid. The oxides N0 2 and 
N 2 3 are themselves readily reduced to nitric oxide. 

If the element upon which the acid acts has not the 
power to take the place of the hydrogen, the action con- 
sists in oxidation. This is shown in the action of strong 
nitric acid upon tin, phosphorus, carbon, sulphur, etc. 
In each case the highest oxidation-product is formed. 
Tin is converted into normal stannic acid, Sn(OH) 4 ; phos- 



282 COLLEGE CHEMISTRY. 

phorus into phosphoric acid, PO(OH) 3 ; carbon into carbon 
dioxide, C0 2 ; and sulphur into sulphuric acid, S0 2 (OH) 2 . 
It disintegrates carbon compounds very readily, converting 
them into their final products of oxidation. In contact 
with the skin it causes bad wounds. 

The acid mostly used in the laboratory has the specific 
gravity 1.2 and contains 32 per cent nitric acid, HN0 3 . 
The commercial acid contains about 68 per cent of the 
acid. 

When a mixture of nitric acid and water is boiled under 
the ordinary atmospheric pressure it loses either water or 
nitric acid until it contains 68 per cent of the acid and 
then it distils over. This does not correspond to any 
definite hydrate of nitric acid, though it approximates the 
composition required by normal nitric acid, N(OH) 5 , or 
HN0 3 -f- 2H 2 0, and it is probable that this hydrate is the 
chief constituent of the mixture. 

Nitric acid is a strong monobasic acid, forming salts of 
the general formula MN0 3 , all of which are soluble in 
water. Because nitric acid is a strong acid, and all 
normal nitrates are soluble in water, it is one of the best 
solvents. 

Red Fuming Nitric Acid is formed in the manufacture 
of nitric acid from saltpetre and sulphuric acid if the tem- 
perature is raised to a sufficient extent to cause the acid 
sulphate to act upon the nitrate : 

NaN0 3 + HNaSO, = Na 2 S0 4 + HN0 3 . 

At this temperature the nitric acid undergoes considerable 
decomposition. The nitrogen peroxide formed is absorbed 
by the nitric acid, and the product thus obtained is the 
red fuming acid. It acts more energetically than nitric 
acid, and finds some applications in the laboratory and in 
the arts. When heated it gives off nitrogen peroxide, and 
if diluted with water it is changed to ordinary nitric acid, 
as nitrogen peroxide is decomposed by water, forming 
nitric acid and nitric oxide or nitrous acid, according to 
the temperature of the water. 



NITROUS ACID. 283 

Nitro-hydrocliloric Acid or Aqua Regia is a liquid formed 
by mixing concentrated nitric and hydrochloric acids. It 
was called aqua regia because it can dissolve gold, the king 
of the metals. The active power of this liquid as a solvent 
of metallic substances is due to the fact that it gives off 
chlorine, and a compound of nitrogen, oxygen, and chlo- 
rine which readily gives up its chlorine. This compound 
has the composition represented by the formula NO CI, and 
it is best designated by the name nitrosyl chloride. The 
product of the action of nitro-hydrochloric acid upon a 
metal is the corresponding chloride. 

Nitrous Acid, HN0 2 . — When certain salts of nitric acid 
are reduced they yield the corresponding nitrites. Thus, 
when potassium nitrate is heated with metallic lead this 
reaction takes place : 

KNO3 + Pb = KN0 2 + PbO. 

Indeed, if potassium nitrate is heated alone it loses oxygen 
and is converted into the nitrite : 

2KN0 3 = 2KN0 2 + 2 ; 

but the reaction is not complete, and the salt thus obtained 
always contains more or less nitrate. 

Nitrous acid is known only in solution. If an attempt 
is made to isolate it from a nitrite the product is the anhy- 
dride, nitrogen trioxide, N 2 3 . Thus, if sulphuric acid is 
added to potassium nitrite the following reaction takes 
place : 

2KN0 2 + H 2 S0 4 = N 2 3 + H 2 + K 2 S0 4 . 

It may be that the first action is the liberation of nitrous 
acid, and that this then breaks down by loss of water. 
The two reactions are represented thus : 

2KN0 2 + H,S0 4 = 2HN0 2 + K 2 S0 4 ; 

2HN0 2 = N 2 3 + H 2 0. 
The trioxide, N 2 3 , however, breaks down into the two 
gases nitrogen tetroxide, N0 2 , and nitric oxide, NO. 



284 COLLEGE CHEMISTRY. 

A certain analogy will be observed between this action 
and that which takes place in the action of potassium 
hydroxide upon an ammonium salt, when ammonium 
hydroxide is probably first given off, and then breaks down 
into ammonia and water. 

Salts are known which are derived from normal nitrous 
acid, N(OH) 3 , but most of the nitrites are derived from the 
acid of the formula NO (OH), which is to be regarded as 
formed from the normal acid by loss of one molecule of 
water : 

N(OH) s = NO(OH) + H 2 0. 

Hyponitrous Acid, H 2 N 2 2 . — The sodium salt of this 
acid is made by reducing sodium nitrite in solution by 
means of sodium amalgam : 

2NaNO s + 8H = Na 2 N 2 2 + 4H 2 0. 

The acid can also be made by oxidation of hydroxylamine. 
It is a solid consisting of white crystalline plates. It is 
very explosive when freed from water. In water solution 
it is much more stable, but at the ordinary temperature it 
breaks down gradually, the principal products being nitrous 
oxide and water: 

H 2 N 2 2 = N 2 + II 2 0. 

Nitrous Oxide, N.,0. — This compound can be obtained 
by reduction of nitric acid, and is sometimes formed in 
considerable quantity when copper is treated with the 
concentrated acid, though when made in this way it is 
always mixed with a large proportion of nitric oxide. The 
best way to make it is to heat ammonium nitrate, 
NI1 4 N0 3 , which breaks down iuto nitrous oxide and water: 

NH 4 NO s = N 2 -f 2II 2 0. 

In the same way we have seen that ammonium nitrite 
breaks down into free nitrogen and water when heated : 

NH 4 NO a = N 2 + 2H 2 0. 



NITROUS OXIDE-NITRIC OXIDE. 285 

In these reactions we see exhibited the tendency of 
hydrogen and oxygen to combine at elevated temperatures. 
At ordinary temperature this tendency is not strong 
enough to cause a disturbance of the equilibrium of the 
parts of the compound. As the temperature is raised and 
the equilibrium thus disturbed, the affinity of the hydrogen 
for the oxygen asserts itself. The two elements combine 
to form water, and the decomposition above represented 
takes place. 

Nitrons oxide is a colorless, transparent gas which has a 
sweetish taste and odor. Its specific gravity is 1.53. It 
is somewhat soluble in water; one volume of water at 0° 
dissolving somewhat more than its own volume of the gas. 
It supports combustion almost as well as pure oxygen. 
Some substances which burn in oxygen do not, however, 
burn in nitrous oxide. Sulphur which burns in oxygen is 
extinguished in nitrous oxide, unless it is previously heated 
to a high temperature. 

When inhaled, nitrous oxide causes a kind of intoxica- 
tion, w T hich is apt to show T itself in the form of hysterical 
laughing. Hence the gas is called laughing gas. Inhaled 
in larger quantity it causes unconsciousness and insensi- 
bility to pain. It is therefore used extensively to prevent 
pain in some surgical operations, particularly in extracting 
teeth. 

When subjected to a low temperature and high pressure 
the gas is easily liquefied, and enclosed in properly con- 
structed metallic cylinders the liquid is now sent into the 
market. In order to get the gas it is only necessary to 
open the stop-cock of the cylinder. When the liquid 
comes in contact with the air it rapidly turns to gas, and 
the temperature is very much lowered in consequence. 
This causes a part of the liquid to solidify. 

Nitric Oxide, NO. — This is the most stable compound 
of nitrogen and oxygen, and is the most common product 
of the reduction of nitric acid. Thus, when nitric acid 
acts upon copper and other metallic elements the chief 
product is generally nitric oxide, though, as we have seen, 



286 COLLEGE CHEMISTRY. 

the reduction may be carried farther. The principal action 
in the case of copper is represented thus : 

8HNO3 + 3Cu = 3Cu(N0 3 ) 2 + 2NO + 4H 2 0. 

Considering the ease with which nitric acid gives up its 
oxygen, and the ease with which copper takes up oxygen, 
it is probable that the copper abstracts oxygen directly 
from the acid as represented thus : 

2HN0 8 + 3Cu = CuO + H 2 + 2NO. 

In this case the copper oxide would at once form copper 
nitrate with the excess of nitric acid : 

6HNO3 + 3CuO = 3Cu(N0 3 ) 2 + 3H 2 0. 

Or, combining the two equations, the total action is repre- 
sented in the same way as it is above. The nitric acid 
must not have a specific gravity higher than 1.2, and the 
temperature must be kept down, otherwise the reduction 
of the nitric acid is carried farther and considerable nitrous 
oxide is formed. 

It is possible that to some extent the hydrogen liberated 
from the acid may act as a reducing agent, thus causing 
the formation of the lower oxides of nitrogen, as, for 
example, 

2HNO3 + 6H = 2NO + 4H 2 0. 

Nitric oxide is a colorless, transparent gas. Its most 
remarkable property is its power to combine directly with 
oxygen when the two are brought together. The act of 
combination is not accompanied by the appearance of 
light, though heat is evolved. In the reaction which takes 
place at ordinary temperatures nitrogen peroxide, N0 2 , is 
formed : 

NO + = N0 2 . 

The product is a colored gas, and the change of the color- 
less nitric oxide to this colored product can therefore easily 



NITROGEN TRIOXIDE— NITROGEN PEROXIDE. 2S7 

be recognized. This reaction is, further, the chief cause 
of the reddish-brown fumes seen when nitric acid acts 
upon metals and other elements. At a low temperature 
some nitrogen trioxide is formed when oxygen acts upon 
nitric oxide. 

From what has already been said, it will appear that in 
nitric oxide the oxygen and nitrogen are more firmly united 
than in the other oxides. Most burning substances are 
extinguished when introduced into it, though a few when 
heated in it to a high temperature extract all or a part of 
the oxygen. Zinc and iron extract half the oxygen and 
convert nitric oxide into nitrous oxide. Potassium and 
sodium decompose it, leaving the nitrogen free. 

Nitrogen Trioxide, N" 2 3 . — This oxide is formed by 
addition of oxygen to nitric oxide at low temperatures; by 
decomposition of the nitrites by means of acids; and by 
the combination of nitric oxide with the peroxide at a 
temperature below — 21°. The gas given off when nitric 
acid is reduced with starch or arsenious oxide, As 2 3 , 
appears to be a mixture of nitric oxide and the peroxide. 
Pure nitrogen trioxide is a liquid of an indigo-blue color. 
At a temperature below 0° it undergoes partial decomposi- 
tion into nitrogen peroxide and nitric oxide : 

N 2 3 = NO + N0 2 . 

With cold water nitrogen trioxide undergoes decomposition 
accompanied by an evolution of nitric oxide. Possibly 
this reaction takes placfi : 

3N 2 3 + H 2 = 2HN0 3 + 4NO. 

By treating the oxide with a solution of sodium hydroxide 
or potassium hydroxide the corresponding nitrite is 
formed : 

2KOH + N 2 3 = 2KN0 2 + H 2 0. 

Nitrogen Peroxide, N0 2 . — "When nitric oxide and oxygen 
are brought together in the proportion of 2 volumes of the 



288 COLLEGE CHEMISTRY. 

former to 1 volume of the latter they combine completely 
to form nitrogen peroxide. These relations will be readily 
understood when it is borne in mind that 2 molecules of 
nitric oxide require 1 molecule of oxygen to effect the 
change, as is shown in the equation 

2NO + 2 = 2N0 2 . 

The compound is most easily obtained by heating lead 
nitrate, when nitrogen peroxide and oxygen are given off, 
and lead oxide remains behind in the vessel : 

Pb(N0 3 ) 2 = PbO + 2NO, + 0. 

If the gases are passed through a tube surrounded by a 
freezing mixture the peroxide is condensed to the form of 
liquid, while the oxygen passes on. When perfectly dry 
the peroxide is easily solidified. It acts energetically upon 
compounds that have the power to take up oxygen. When 
treated with water it undergoes decomposition. If the 
temperature is low, nitrous and nitric acids are formed : 

2N0 2 + H 2 = HN0 2 + HN0 3 . 

If the water is hot, however, the products are nitric acid 
and nitric oxide: 

3N0 2 + H 2 = 2HN0 3 + NO. 

The nitric oxide thus formed will take up oxygen from the 
air and yield nitrogen peroxide again, and this, in contact 
with hot water, will be decomposed, forming nitric acid 
and nitric oxide, until all the peroxide is converted into 
nitric acid. 

The determinations of the specific gravity of the gas 
from the peroxide show that at low temperatures the molec- 
ular formula is N 2 4 , but that when the temperature 150° 
is reached the molecule is represented by the formula N0 2 . 
The compound appears therefore to undergo gradual 



COMPOUNDS OF NITROGEN. 289 

decomposition or dissociation by heat, so that until the 
temperature 150° is reached the gas is a mixture of the 
compounds N 2 4 and N0 2 . 

Nitrogen Pentoxide, N 2 5 . — This compound, which 
bears to nitric acid the relation of an anhydride, is formed 
by passing chlorine over silver nitrate and condensing the 
product. The reaction takes place thus: 

2AgNO s + Cl 2 = N 2 5 + 2AgCl + 0. 

It is also formed by treating nitric acid with phosphorus 
pentoxide, P 2 5 , a compound that has a very marked 
power to unite with water. The action is represented 
thus : 

2HN0 3 = TSfi s + H 2 0. 

The pentoxide is a crystallized substance, which readily 
decomposes into nitrogen peroxide and oxygen. In conse- 
quence of the ease with which it gives up its oxygen it acts 
violently upon many oxidizable substances. With water it 
forms nitric acid: 

N 2 5 + H 2 = 2IINO3. 

Compounds of Nitrogen with the Elements of the 
Chlorine Group. — Notwithstanding the ease with which 
chlorine combines with most elements, and the stability of 
the compounds which it forms with them, its compound 
with nitrogen is extremely unstable. It can be made by 
the action of chlorine on ammonia, and by decomposing a 
solution of ammonium chloride by means of an electric 
current. In the latter case chlorine is liberated at one of 
the poles and then acts upon the ammonium chloride : 

NH 4 C1 + 6C1 = 4HC1 + NOI3. 

It appears that when chlorine acts upon ammonia different 
products are formed by the substitution of chlorine for the 
hydrogen atom for atom, thus : 



290 COLLEGE CHEMISTRY. 

NH 8 + Cl 2 = NH a Cl + HCl; 
NH 2 C1 + Cl 2 = NHC1 2 + HOI; 
NHC1 2 + Cl 2 = NCI, + HC1. 

According to this, the trichloride of nitrogen is the final 
product of the substituting action of chlorine upon am- 
monia. The compound is an oil, which undergoes decom- 
position very readily. It is, indeed, one of the most 
explosive substances known. It is decomposed by heat, 
and especially by contact with certain substances, among 
which are oil of turpentine and caoutchouc. It is slowly 
decomposed by water, though, probably owing to the slight 
affinity of nitrogen for oxygen, the decomposition does not 
take place as readily as that of the compounds of sulphur 
and chlorine. Direct sunlight causes explosion of the 
chloride. 

When ammonia is treated with iodine reactions take 
place similar to those which take place with chlorine. 
The product is the iodide of nitrogen, N 2 H 3 I 3 . This com- 
pound,' like the chlorine compounds, is extremely explo- 
sive. The simplest way to prepare it is to place a little 
powdered iodine on a filter and pour concentrated ammonia 
over it. The substance should be made in only very small 
quantities at a time. When dried it decomposes with 
violent explosion by contact even with soft substances. 

EXPERIMENTS. 
Preparation and Properties of Ammonia. 

Experiment 124. — To a little ammonium chloride on a watch- 
glass add a few drops of a strong solution of caustic soda, and 
notice the odor of the gas given off. Do the same thing with 
caustic potash. Mix small quantities pf ammonium chloride and 
lime in a mortar, and add a few drops of water. 

Experiment 125. — Mix 20 parts iron filings, 1 part potassium 
nitrate, and 1 part solid potassium hydroxide, and heat the mix- 
ture in a test-tube. Is there any evidence of the formation of 
ammonia? 

Experiment 126. — Arrange an apparatus as shown in Fig. 37, 
omitting, however, the funnel-tube ; a cork with one hole will 



EXPERIMENTS WITH AMMONIA. 



291 



therefore serve. Weigh 100 grains quicklime in the flask, and 

add just enough water to slake it without making it moist ; then 

add 50 grams ammonium chloride, and mix by shaking. Heat on 

a sand-bath. After the air is driven out, the 

gas will be completely absorbed by the water 

in the first Woulff's flask if shaken from time 

to time. Disconnect the delivery-tube from 

the series of Woulff's flasks, and connect with 

another tube bent upward. Collect some of 

the gas by displacement of air, placing the 

vessel with the mouth downward. (Why ?) 

The arrangement is shown in Fig. 52. The 

vessel in which the gas is collected should be 

dry, as water absorbs ammonia very readily. 

Hence, also, it cannot be collected over water. 

In the gas collected introduce a burning 

stick or taper. Ammonia does not burn in 

air, nor does it support combustion. In 

working with the gas great care must be taken 

to avoid inhaling it in any quantity. After 

enough has been collected in cylinders to 

exhibit the chief properties, connect the delivery-tube again with 

the series of Woulff's flasks, and pass the gas over the water as 

long as it is evolved. 

Ammonia Burns in Oxygen. 

Experiment 127. — Put a little of a concentrated solution of 
ammonia in a flask placed upon a tripod. Heat gently and, from 
a gasometer, pass a rapid current of oxygen through a bent tube 
into the liquid. Apply a light to the mouth of the vessel, when 
the ammonia will be seen to burn. 




Fig. 52. 



Ammonia forms Ammonium Salts with Acids. 

Experiment 128. — Put 100 cc. of a dilute solution of ammonia 
in an evaporating-dish. Try its effect on red litmus paper. Slowly 
add dilute hydrochloric acid until the alkaline reaction is destroyed 
and the solution is neutral. Evaporate to dryness on a water- 
bath. Compare the substance thus obtained with sal-ammoniac, 
or ammonium chloride. Taste. Heat on a piece of platinum- 
foil. Treat with a caustic alkali. Treat with a little concentrated 
sulphuric acid in dry test-tubes. Do they appear to be identical ? 
Similarly sulphuric acid and ammonia yield ammonium sulphate; 
nitric acid and ammonia yield ammonium nitrate ; etc. 



292 



COLLEGE CHEMISTRY. 



Experiment 129. — Fill a dry cylinder with ammonia gas, and 
another of the same size with hydrochloric acid gas. Bring them 
together with their mouths covered. Quickly remove the covers, 
when a dense white cloud will appear in and about the cylinders. 
This will soon settle on the walls of the vessels as a light white 
solid. It is ammonium chloride. Thus, from two colorless gases 
we get a solid substance by an act of chemical combination. Heat 
is evolved in the act of combination. 



Composition of Ammonia. 

Experiment 130.— This experiment should be performed by a 
person experienced in the use of chemical apparatus. A glass 
tube, such as represented in Fig. 53, provided with a glass stop- 
cock is needed. Fill this tube 
with chlorine free from air 
over a saturated solution of 
sodium chloride. After it is 
filled let it stand for some 
time mouth downward in the 
solution of sodium chloride 
to let the liquid drip out of 
it. Close the stop-cock and 
remove the tube from the 
solution. Hold it mouth up- 
ward, and pour a concen- 
trated solution of ammonia 
into the funnel-like projec- 
tion above the stop-cock; put 
in the glass stopper, and now 
by slightly opening the stop- 
cock let the ammonia pass 
drop by drop into the tube. 
Keaction between the chlo- 
rine and the ammonia takes 
place, accompanied by a 
marked evolution of heafe, 
and in a partly-darkened 
room light is seen. Great 
care must be taken not to 
admit air with the ammonia. 
After nearly all the ammonia 
has passed in from the fun- 
nel, pour into the funnel about two-thirds as much ammonia as 




EXPERIMENTS WITH NITRIC ACID. 



2 93 



has already been used, and let this in gradually. Leave the stop- 
cock closed, and fill the funnel with dilute sulphuric acid. Fit a 
bent tube into a cork; fill this tube with dilute sulphuric acid ; 
put the cork in the funnel, and the other end of the tube in a 
small beaker containing dilute sulphuric acid, and, after immers- 
ing the long tube in water of the ordinary temperature, open the 
stop-cock. If the operation has been carried out as it should 
be, the dilute acid will flow into the tube until it is two-thirds 
full, and will then stop. The residual gas is nitrogen. What 
evidence in regard to the composition of ammonia is furnished 
by this experiment ? 

The arrangement of the apparatus in the last stage of the ex- 
periment is shown in Fig. 53. 

Preparation and Properties of Nitric Acid. 

Experiment 131.— Arrange an apparatus as shown in Fig. 54. 
In the retort put 20 grams sodium nitrate (Chili saltpetre) and 20 




Pig. 54. 

grams concentrated sulphuric acid. On gently heating, nitric 
aeid will distil over, and be condensed in the receiver. — What re- 
action takes place ? — After the acid is all distilled off, remove the 
contents of the retort. Recrystallize the substance from water, 
and compare it with the sodium sulphate obtained in the prepara- 
tion of hydrochloric acid. (See Experiment 74.) In the latter stage 
of the operation the vessels become filled with a reddish-brown 
gas. The acid which is collected has a somewhat yellowish color. 
Experiment 132.— Mix together 400 grams concentrated sul- 
phuric acid and 80 grams ordinary concentrated nitric acid. 



2 9 4 



COLLEGE CHE MIS TR Y. 



Pour the sulphuric acid into the nitric acid. Distil the mixture 
from a retort arranged as in the preceding experiment, taking 
care to keep the neck of the retort cool by placing filter-paper 
moistened with cold water on it. Use the acid thus obtained for 
the purpose of studying the properties of pure nitric acid. 

Xitric Acid gives up Oxygen readily, axd is hence 
a good Oxidizing Agent. 

Experiment 183. — Pour concentrated nitric acid into a wide 
test-tube, so that it is about one-fourth filled. Heat the end of a 
stick of charcoal of proper size, and, holding the other end with 
a forceps, introduce the heated end into the acid. It will continue 
to burn with a bright light, even though it is placed below the 




surface of the liquid. The action is oxidation. The charcoal in 
this case finds the oxygen in the acid and not in the air. Great 
care must be taken in performing this experiment. The charcoal 
should not come in contact with the sides of the test-tube. A large 
beaker-glass should be placed beneath the test-tube, so that 
in case it breaks the acid will be caught and prevented from doing 
harm. The arrangement of the apparatus is shown in Fig. 55. 

The gases given off from the tube are offensive and poisonous. 
Hence this experiment as well as all others with nitric acid should 
be carried on under a hood in which the draught is good. 

Experiment 134.— Boil a little strong nitric acid in a test-tube 
in the upper part of which some horse-hair (or woollen yarn) has 
been introduced in the form of a stopper. The horse hair (or yarn) 
will take fire and burn, and leave a white residue. Hold the test- 



EXPERIMENTS WITH NITRIC ACID. 295 

tube with a forceps over a vessel to catch the contents should the 
tube break. 

Experiment 135. — In a small flask put a few pieces of granu- 
lated tin. Pour on this just enough strong nitric acid to cover it. 
Heat gently over a small flame. Soon action will take place. Col- 
ored gases will be evolved, the tin will disappear, and in its place 
will be found a white powder. This consists mostly of tin and 
oxygen. (See Experiment 13. 

Metals dissolve in Nitric Acid, forming Nitrates. 

Experiment 136. — Dissolve a few pieces of copper-foil in or- 
dinary commercial nitric acid diluted with about half its volume 
of water. The operation should be carried on in a good-sized 
flask and under an efficient hood. When the copper has disap- 
peared, pour the blue solution into an evaporating- dish, and 
evaporate down to crystallization. Compare the substance thus 
obtained with copper nitrate. Heat specimens of each. Treat 
small specimens with sulphuric acid. What evidence have you 
that the two substances are identical ? 

Nitrates are decomposed by Heat. 

Experiment 137. — Heat some potassium nitrate in a test-tube. 
Introduce a piece of wood with a spark on it. Heat also lead 
nitrate, copper nitrate, and any other nitrates that may be 
available. What difference do you observe between the decom- 
position of potassium nitrate and that of lead nitrate ? 

Nitrates are Soluble in Water. 

Experiment 138. — Try the solubility in water of the nitrates 
used in the last experiment. 

Nitric Acid is reduced to Ammonia by Nascent 
Hydrogen. 

Experiment 139.— In a good-sized test-tube treat a few pieces 
of granulated zinc with dilute sulphuric acid. What is evolved ? 
Prove it. Now add drop by drop dilute nitric acid. The hydro- 
gen ceases to be given off. Pour the contents of the tube into an 
evaporating-dish and evaporate the liquid. Put the residue into 
a test-tube and add a solution of caustic soda, when the smell of 
ammonia will be noticed. Try the action of the gas on red litmus 
paper. Moisten the end of a glass rod with a little hydrochloric 
acid, and hold it in the tube. AVhite fumes are seen. What are 
they? Do the same with nitric acid. What are the fumes in 
this case ? 



296 



COLLEGE CHEMISTRY. 



Nitrous Acid. 

Experiment 140. — Melt 25 grams potassium nitrate in a shal- 
low iron plate and gradually add 50 grams metallic lead cut in 
small pieces. Stir them together as thoroughly as possible. After 
the mass is cooled down, break it up and treat with water in a 
flask. The potassium nitrate will dissolve, while the lead oxide 
and unused lead will not dissolve. Filter. Add a little sulphuric 
acid to some of the solution. A colored gas will be given off. See 
whether a solution of potassium nitrate acts in the same way. 
Treat with sulphuric acid a little of the residue left after heating 
potassium nitrate alone in a test-tube as in Experiment 138. 

Nitrous Oxide. 

Experiment 141. — In a retort heat 10 to 15 grams crystallized 
ammonium nitrate until it has the appearance of boiling. Do 
not heat higher than is necessary to secure a regular evolution of 
gas. Connect a wide rubber tube directly with the neck of the 
retort and collect the evolved gas over water, as in the case of 
oxygen. It supports combustion almost as well as pure oxygen. 
Try experiments with wood, a candle, and a piece of phosphorus. 

Nitric Oxide. 

Experiment 142.— Arrange an apparatus as shown in Fig. 56. 
In the flask put a few pieces of copper-foil. 
Cover this with water. Now add slowly, 
waiting each time for the action to begin, 
ordinary concentrated nitric acid. When 
enough nitric acid has been added gas will be 
evolved. If the acid is added rapidly, it not 
infrequently happens that the evolution of 
gas takes place too rapidly, so that the liquid 
is forced out of the flask through the funnel- 
tube. This can be avoided by not being in a 
hurry. At first the vessel becomes filled with 
a reddish-brown gas, but soon the gas evolved 
becomes colorless. Collect over water two or 
three vessels full. The gas collected is prin- 
cipally nitric oxide, NO, though it is fre- 
quently mixed with a considerable quantity 
of nitrous oxide. 

Experiment 143. — Turn one of the vessels 
containing colorless nitric oxide with the fig. 56. 

mouth upward, and uncover it. The colored gas is at once seen, 




EXPERIMENTS IVITH NITRIC OXIDE, ETC 297 

presenting a very striking appearance. Do not inhale the gas. 
Perforin the experiments with nitric oxide where there is a good 
draught. 

Experiment 144.— Pass nitric oxide into a concentrated solu- 
tion of ferrous sulphate. Afterwards heat the solution and col- 
lect the gas. What do you conclude that the gas is ? 

Nitrogen" Trioxide. 

Experiment 145.— In a flask fitted with a safety-funnel and a 
delivery-tube pour nitric acid of specific gravity 1.30-1.85 upon 
coarsely-granulated arsenious oxide, As 2 3 . Heat gently, and 
conduct the gases through a tube surrounded by a freezing mix- 
ture, as in Experiment 110. 

Nitrogen Peroxide. 

Experiment 146.— Admit a little air to nitric oxide contained 
in a bell-jar over water, and let the vessel stand. Almost imme- 
diately the color will disappear, showing that the nitrogen perox- 
ide formed is decomposed. Again admit air, and let the vessel 
stand. The same changes will be noticed as in the first instance. 
If oxygen is used instead of air the above changes can be repeated 
over and over again. Devise an experiment for the purpose of 
determining whether the nitric oxide is gradually used up or not. 



CHAPTER XVII. 

ELEMENTS OF FAMILY V, GROUP B: 

PHOSPHORUS.— ARSENIC— ANTIMONY.— BISMUTH. 

THE ELEMENTS AND THEIR COMPOUNDS WITH 

HYDROGEN. 

General. — The elements of this group bear to nitrogen 
much the same relations that the members of the sulphur 
group bear to oxygen, and those of the chlorine group bear 
to fluorine. In general they form compounds of the srme 
character and of similar composition. At the same time 
gradations in properties are noticed in passing from one 
end of the group to the other. Like nitrogen, the 
elements of the group are strongly marked acid-formers, 
though this character grows less marked from nitrogen to 
bismuth. Antimony is both an acid-forming and a base- 
forming element, while bismuth is more basic than acid. 
The stability of the hydrogen compounds decreases from 
nitrogen to antimony; while bismuth does not form a 
compound with hydrogen. Ammonia, as we have seen, is 
strongly basic ; the corresponding compound of phosphorus 
and hydrogen has weak basic properties, while those of 
arsenic and antimony have no basic properties. These 
hydrogen compounds correspond in composition to am- 
monia. They are: 

NH 3 : PH 3 AsH 3 SbH 3 

With chlorine they all form compounds corresponding to 
nitrogen trichloride, and phosphorus and antimony form 
compounds in which they are quinquivalent, while bismuth 

298 



ELEMENTS OF FAMILY V, GROUP B. 299 

forms a chloride, Bi 2 Cl 4 , in addition to the trichloride. 
The compounds referred to are: 

BiA 

NC1 3 : PCI3 As01 3 SbCl 3 BiCl 3 
PCI. SbCl 5 

They all form two oxides corresponding to nitrogen tri- 
oxide and pentoxide: 

N 2 8 : P 2 3 As 2 3 Sb.,0 3 Bi 9 3 
N 2 5 : P 2 5 As 2 5 Sb 2 5 Bi 2 5 

None of the elements of the group forms as great a variety 
of compounds with oxygen as nitrogen does. Antimony, 
however, forms the oxide Sb 2 4 , corresponding to nitrogen 
peroxide, N 2 4 ; and bismuth forms the oxide Bi 2 2 or 
BiO, corresponding to nitric oxide, NO. 

The hydroxyl compounds or acids, like those of nitro- 
gen, are related to the maximum hydroxyl compounds of 
the elements with the valence 5, and to the maximum 
hydroxyl compounds of the elements with the valence 3. 
That is to say, they may be regarded as derived from a 
hydroxide of the general formula M(OH)_, and another of 
the formula M(OH) 3 . Where M is nitrogen these acids 
break down to the forms N0 2 (OH) and NO(OH) by loss 
of one or two molecules of water. In the case of the ele- 
ments of the phosphorus group, however, the breaking 
down is not generally carried as far as with nitrogen. The 
general rule is the same as in the sulphur and chlorine 
groups : the normal acid breaks down to form compounds 
containing the same number of hydrogen atoms as the 
hydrogen compounds of the elements. Thus the hydroxyl 
derivatives of chlorine generally break down to form com- 
pounds containing one atom of hydrogen, or the same 
number that is contained in the hydrogen compound, hy- 
drochloric acid, thus: Cl(OH) r yields C10 3 (OH); Cl(OH). 
yields C10 2 (OH), etc. So, also, in the sulphur group, 
S(OH) 6 yields S0 2 (OH) a , etc., the number of hydrogen 



3°o COLLEGE CHEMISTRY. 

atoms in the common form of the acid being the same as 
that in the hydrogen compound of sulphur, SH 2 . 

Of the elements of this group phosphorus occurs most 
abundantly in nature, arsenic and antimony next, and 
bismuth least abundantly. Arsenic, antimony, and bis- 
muth occur to some extent in the uncombined condition. 
Phosphorus occurs in combination. All the elements of 
the group find applications in the arts, either as the ele- 
ments or in the form of compounds. 

Phosphorus, P (At. Wt. 31). 

Occurrence. — The name phosphorus is derived from the 
Greek 0g3s", light, and cpopos, carrier, on account of the 
fact that it gives light and takes fire very easily. The 
element occurs in nature in the form of phosphates derived 
from orthophosphoric acid, H 3 P0 4 . The chief of these 
is calcium phosphate, Ca 3 (POJ 2 , which is the principal 
constituent of the minerals phosphorite and apatite, and 
of the ashes of bones. The phosphates, like the nitrates, 
are widely distributed in the soil and are of fundamental 
importance in the process of plant life. The phosphates 
found in the bones are taken into the animal body in the 
food. All plants used as food contain small quantities of 
the phosphates which they get from the soil. The phos- 
phates taken into the body are partly given off in the 
excrement and urine, and it was in an examination of 
urine made in the hope of finding the philosopher's stone 
that phosphorus was discovered in 1669. At present phos- 
phorus is made almost entirely from bones. 

Preparation. — Besides the phosphates, considerable 
quantities of organic materials are contained in bones. 
When the bones are burned the organic materials pass off 
for the most part in the form of carbon dioxide, water, 
and volatile compounds containing nitrogen, and the 
so-called mineral or earthy portions, the chief constituent 
of which is tertiary calcium phosphate, Ca 3 (POJ 2 , or 
phosphoric acid in which calcium has been substituted for 



PHOSPHORUS. 3 OT 

all the hydrogen, remain behind. The tertiary phosphate 
is insoluble in water, and there is no simple way by which 
the phosphorus can be set free from it. When it is 
treated with sulphuric acid calcium sulphate which is 
difficultly soluble is deposited and mono-calcium phosphate, 
Ca(H 2 P0 4 ) 2 , is in the solution. The reaction is repre- 
sented as follows: 

Ca 3 (P0 4 ) 2 + 2H 2 S0 4 = Ca(H 2 P0 4 ) 2 + 2CaS0 4 . 

The calcium sulphate, or gypsum, is allowed to settle 
and is then filtered off and washed. The solution is 
evaporated to dryness and heated. The mono-calcium 
phosphate is by this means converted into calcium meta- 
phosphate : 

Ca(H,PO ( ) 2 = Ca(P0 3 ) 2 + 2H 2 0. 

This metaphosphate is then mixed with charcoal and 
generally with sand, when the reaction represented by the 
following equation takes place: 

2Ca(P0 3 ) 2 + 2Si0 2 + IOC = 2CaSi0 3 + 10CO + 4P. 

If an electric furnace is used for the heating it is not 
necessary to convert the ordinary or tertiary phosphate into 
the mono-calcium phosphate, as the former can be reduced 
directly by charcoal and sand : 

2Ca 3 (P0 4 ) 2 + 6Si0 2 + IOC = 3CaSi0 3 + 10CO + 4P. 

The phosphorus passes over in the form of vapor, and is 
collected under water. The crude phosphorus thus 
obtained must be subjected to a cleansing process before it 
can be used. For this purpose it is pressed, while in the 
molten condition under water, through chamois leather, 
or it is distilled again from iron retorts; or, still better, it 
is treated with chromic acid as follows : It is fused under 
water, then a little potassium or sodium bichromate in 
solution is added, and afterwards an equivalent proportion 



302 COLLEGE CHEMISTRY. 

of sulphuric acid, and the whole allowed to stand for two 
hours or more. The phosphorus is then washed with hot 
water, and after being siphoned off it is filtered through 
canvas bags. The phosphorus is then cast into sticks in 
tin tubes. In this form it generally comes into the market. 

At the time of the last report available there were manu- 
factured in one year about 3000 tons of phosphorus in 
England and in France. Quite recently phosphorus has 
been manufactured to some extent in Sweden. 

Properties. — Ordinary phosphorus is colorless or slightly 
yellowish, translucent, and at ordinary temperatures it can 
be cut like wax, but it becomes hard and brittle at low 
temperatures. It melts at 44°, and boils at 290°. It is 
insoluble in water. When kept under water for some time 
in dispersed light it becomes opaque, crystalline on the 
surface, and yellow. It is soluble in carbon disulphide, 
and crystallizes when deposited from this solution. It 
gives off fumes in contact with the air, and emits a pale 
light which is known as a phosphorescent light. It is very 
poisonous, the inhalation of the vapor in small quantities 
causing very serious disturbance of the system. The 
workmen in the factories where phosphorus is made or 
used are frequently affected by phosphorus-poisoning. 
Among the prominent symptoms is gradual decomposition 
of the bones. When taken into the stomach phosphorus 
also acts as a poison and causes death. When heated in 
the air it takes fire at 50°. It also takes fire by rubbing, 
and it must be handled with the greatest care, as wounds 
caused by it are dangerous and difficult to heal. When it 
burns in the air it is converted into the pentoxide, P 2 5 , 
which is also the product of its combustion in oxygen, as 
we have seen. It combines also with other elements 
directly, frequently with evolution of light. Thus, when 
it is brought together with chlorine, bromine, and iodine, 
it forms the compounds PC1 3 , PBr 3 , and PI 3 . It also 
combines with sulphur. When a piece is put in water and 
the water boiled, a part of the phosphorus passes over, 
and if the water-vapor is condensed in a glass tube in a 



PHOSPHORUS. 3°3 

dark room, it is seen to be phosphorescent. This fur- 
nishes a convenient method for its detection, as, for 
example, in a case of suspected poisoning by phosphorus. 

When phosphorus is left for a long time under water in 
the light, it becomes at first yellow, then reddish, and 
finally red. The same change takes place when it is 
heated for a time in an atmosphere which is free from 
oxygen; and rapidly when it is heated to 300° in an 
hermetically-sealed tube. The red substance thus obtained 
has properties entirely different from those of ordinary 
phosphorus. It is a red powder, that frequently has a 
crystalline structure. It does not emit light. It does not 
melt at a low temperature. It is not poisonous, and can- 
not be easily ignited. Further, it is perfectly insoluble in 
carbon disulphide. In every respect this red modification 
of phosphorus conducts itself as a much less active sub- 
stance chemically than ordinary phosphorus. In an 
atmosphere of carbon dioxide it is converted into ordinary 
phosphorus when heated to 261°, and if heated to this 
temperature in the air it takes fire, and then forms the 
same product that ordinary phosphorus does in burning. 

Treated with oxidizing agents, as, for example, nitric 
acid, phosphorus is slowly converted into phosphoric acid ? 
just as sulphur is converted into sulphuric acid under the 
same conditions. 

Applications of Phosphorus. — Phosphorus is used prin- 
cipally in the manufacture of matches and as a poison for 
vermin. Various mixtures are used for making matches. 
Nearly all of them contain phosphorus together with some 
oxidizing compound, and some neutral substance to act as 
a medium for holding the constituents together. An 
example is a mixture consisting of 2 parts phosphorus 
1 part manganese dioxide, 3 parts chalk, -J part lamp- 
black, and 5 parts glue. The mixture used in the manu- 
facture of the so-called "safety matches" consists of 
potassium chlorate, potassium bichromate, minium, and 
antimony trisulphide. This will not ignite by simple fric- 
tion, but will ignite when drawn across a paper upon 



3°4 COLLEGE CHEMISTRY. 

which is a mixture of red phosphorus and antimony penta- 
sulphide. 

Compounds of Phosphorus with Hydrogen There are 

three compounds of phosphorus with hydrogen, a gaseous 
compound of the formula PH 3 , corresponding to am- 
monia; a liquid of the formula PH 2 , or P 2 H 4 , correspond- 
ing to hydrazine; and a solid of the formula P 2 H, or P 4 H 2 . 

Phosphine, Gaseous Phosphuretted Hydrogen, PH 3 . — 
This compound is formed: 

(1) By treating a strong solution of potassium hydroxide 
with phosphorus, when reaction takes place as follows : 

3KOH + 4P + 3H 2 = 3KH 2 P0 2 + PH 3 . 

The compound KH 2 P0 2 is known as potassium hypophos- 
phite, being derived from hypophosphorous acid, H 3 P0 2 . 

(2) By treating zinc phosphide with dilute hydrochloric 
acid. Assuming that zinc phosphide has the composition 
represented by the formula Zn 3 P 2 , the reaction with 
hydrochloric acid takes place according to the equation 

Zn 3 P 2 + 6HC1 = 3ZnCl 2 + 2PH 3 . 

(3) By treating phosphonium iodide, PH 4 I, with water 
or a dilute solution of potassium hydroxide: 

PII 4 I + H 2 = PH 3 + HI + H 2 0; 
PH 4 I + KOH = PH 3 + KI + H 2 0. 

When made from phosphorus and potassium hydroxide it 
always contains a considerable proportion of hydrogen, for 
the reason that potassium hypophosphite gives off hydrogen 
when heated with a solution of potassium hydroxide. 
From zinc phosphide and from phosphonium iodide it can 
be obtained in pure condition. 

Phosphine is a colorless gas with an unpleasant, garlic- 
like odor. It is insoluble in water, and is poisonous. It 
burns, but does not take fire spontaneously when pure. 

Although pure phosphine does not take fire spon- 
taneously when brought in contact with the air, the gas 



ARSENIC. 3° 5 

made by any one of the methods above referred to is pretty 
sure to contain some of the liquid compound of phosphorus 
and hydrogen, .P,H 4 , which is spontaneously inflammable, 
and therefore the gas takes fire. If it is collected in a 
glass vessel over water, and allowed to stand so that the 
light acts upon it, the liquid phosphine is decomposed into 
the gaseous and solid varieties, and the residual gas no 
longer has the property of taking fire spontaneously. 

Arsenic, As (At. Wt. 75). 

Occurrence. — Arsenic occurs in nature to some extent 
in the uncombined condition or native. Compounds of 
the metals with arsenic, or the arsenides, occur very widely 
distributed, and they frequently accompany, and are 
similar to, the sulphides. The most common compound 
of this kind is the so-called arsenical pyrites, which has 
the composition FeAsS, and may therefore be regarded as 
iron pyrites, FeS 2 , in which one atom of arsenic has been 
substituted for one atom of sulphur. Among other arsenic 
compounds deserving special mention are the two arsenides 
of iron of the formulas FeAs 2 and Fe 2 As 3 , which are 
apparently analogous to the sulphides FeS 2 and Fe 2 S 3 ; 
and, further, the sulphides of arsenic, orpiment, As 2 S 3 , 
and realgar, As 2 S 2 . The oxide As 2 3 occurs in consider- 
able quantity, and also salts of arsenic acid, or the 
arsenates, which in composition are analogous to the 
phosphates. 

Preparation. — The arsenic which comes into the market 
is either that which occurs native or it is made from 
arsenical pyrites by heating : 

FeAsS = FeS + As. 

Properties. — Arsenic has a metallic lustre and steel 
color. It is very brittle. When heated it volatilizes with- 
out melting. At red heat it burns with a bluish flame, 
and the vapor given off has the odor of garlic. This odor 
produced under such circumstances is very characteristic 



306 COLLEGE CHEMISTRY. 

of arsenic, and furnishes one of the means for detecting 
it. Arsenic combines with most elements directly, the 
action being accompanied in some cases, as in that of 
chlorine, by an evolution of light. As an element it is 
not poisonous, but when oxidized to the form of the oxide 
As 2 3 it is extremely poisonous. As it is easily oxidized, 
the element itself may act as a poison. 

Arsine, Arseniuretted Hydrogen, AsH 3 . — This com- 
pound is analogous to ammonia and to gaseous phosphine. 
It is made by reduction of compounds of arsenic containing 
oxygen, as arsenic trioxide or arsenic acid; and also by 
treating a compound of zinc and arsenic with dilute sul- 
phuric acid. The reactions involved in the first method 
are 

As 2 3 + 6H 2 = 2AsH 3 + 3H 2 0; 

H 3 As0 4 + 4H 2 = AsH 3 + 4H 2 0. 

That involved in the second method mentioned is: 

As 2 Zn 3 + 3H 2 S0 4 = 2AsH 3 + 3ZnS0 4 . 

Arsine is a colorless gas with an odor suggestive of 
garlic. It is extremely poisonous, even very small quan- 
tities being capable of producing bad effects, and it requires 
but little to cause death. When ignited in the air it takes 
fire and burns with a pale blue flame, the products of the 
combustion being arsenic trioxide, As 2 3 , and water. The 
gas is so unstable that, when it is passed through a glass 
tube heated to redness, it is decomposed into arsenic and 
hydrogen, the former being deposited just in front of the 
heated portion of the tube as a thin, almost black, layer 
with a high metallic lustre. 

Antimony, Sb (At. Wt. 120). 

Occurrence. — Antimony occurs in nature in small quan- 
tities in the free state, but chiefly in the form of stibnite, 
which is the trisulphide Sb 2 S 3 . This also occurs very 
widely distributed in nature in combination with sulphides 



ANTIMONY. 307 

of various metals, as copper, lead, and silver. The element 
is made from the sulphide either by heating it with iron, 
with which the sulphur combines, leaving the antimony 
free; or by roasting it, that is, heating it in combination 
with the air, thus converting the antimony into the 
tetroxide Sb 2 4 , and the sulphur into the dioxide S0 2 , 
and then treating the oxide of antimony with reducing 
agents, as, for example, carbon : 



Sb.^0, + 4C = 2Sb -f 4C0. 



Properties. — Antimony is hard and brittle; has a silver- 
white color, and a high metallic lustre. It can be distilled 
at white heat. At ordinary temperature it is not changed 
by contact with the air. When heated to a sufficiently 
high temperature in the air it takes fire and burns, form- 
ing the white oxide Sb 2 3 . It combines directly with 
chlorine, forming the chloride SbCl 5 . Mtric acid oxidizes 
it either to antimony oxide, Sb 2 3 , or antimonic acid, 
H 3 Sb0 4 . Aqua regia dissolves it. Hot concentrated sul- 
phuric acid dissolves it, forming antimony sulphate, and 
sulphur dioxide escapes. 

Applications of Antimony. — Antimony is used as a con- 
stituent of several alloys, which are somewhat indefinite 
compounds which metallic elements form with one another. 
Among the alloys of antimony are type-metal, from which 
type is made, and britannia metal. The former consists 
of lead and antimony, and the latter of tin and antimony. 
There are a number of alloys that contain antimony. 
These will be referred to under the other constituents. 

Stibine, SbH 3 . — This analogue of ammonia, phosphine, 
and arsine is more like arsine than it is like the others. 
It is made by the same methods as those used in making 
arsine, i. e. , by treating an alloy of zinc and antimony with 
sulphuric acid, or by reducing oxides of antimony by means 
of nascent hydrogen. The latter method gives a gas con- 
taining a large percentage of hydrogen, but for most pur- 
poses this is not objectionable. It is only necessary to 
introduce into a flask containing zinc and dilute sulphuric 



308 COLLEGE CHEMISTRY. 

acid a little of a solution of some oxygen compound of 
antimony, when the reduction is at once effected, and the 
escaping hydrogen contains stibine. 

Stibine is a colorless, inodorous gas, which burns with a 
greenish-white flame. In general, it conducts itself much 
like arsine. It is unstable and breaks down when the tube 
through which it is passing is heated to about 150°. It 
then leaves a deposit which looks like that formed in the 
case of arsine. When a cold object, as a piece of porcelain, 
is held for a moment in a flame of stibine a dark deposit 
is formed which resembles that formed with arsine. 

Methods of Distinguishing between Arsenic and Anti- 
mony. — As arsenic is sometimes used in cases of poisoning, 
the question of deciding whether it is present in a given 
liquid or mixture is of great importance. One of the chief 
difficulties encountered is the similarity of the two elements 
arsenic and antimony. The method commonly employed 
in examining a substance for arsenic is known as Marsh's 
test. This consists in getting the substance in solution, 
and then pouring some of the liquid into a vessel contain- 
ing pure zinc and pure dilute sulphuric acid. If arsenic 
is present in the solution it will, under these circum- 
stances, be converted into arsine, the presence of which 
can be recognized by heating the tube through which the 
gas is passing, and by holding a piece of porcelain in the 
flame. If deposits are not formed in the tube or on the 
porcelain, arsenic is not present; but if deposits are 
formed, the only conclusion that can be drawn is that 
either arsenic or antimony is present, or possibly both may 
be present. For the purpose of distinguishing between 
the two elements, advantage is taken of the following 
differences between the spots: The antimony spots are 
darker than those formed by arsenic, and they have a 
smoky appearance, while those of arsenic have not; 
further, the arsenic deposits are quite volatile, and can 
therefore be driven before the flame in the tube or upon 
the porcelain, while those of antimony are not volatile; 
again, the deposits of arsenic are easily soluble in a solu- 



BISMUTH. 309 

tion of sodium hypochlorite or hypobromite, while the 
antimony deposits are insoluble in these solutions. There 
are other differences, but those mentioned will suffice to 
enable a careful worker and observer to distinguish between 
the two without any possibility of doubt. Another diffi- 
culty always encountered in examining for arsenic is the 
fact that the sulphuric acid, the zinc, and the glass of 
which the vessels are made may contain arsenic. It is 
quite possible to overcome all the difficulties and to decide 
positively whether arsenic is present or not. If it is found 
that on heating the tube through which the Irydrogen is 
passing no deposit is formed, even after continued heating, 
and that the hydrogen flame gives no deposit upon a piece 
of porcelain introduced into it, then it is safe to proceed 
with the examination of the suspected liquid. If the sub- 
stance which is to be examined for arsenic has to be treated 
with chemical compounds in order to prepare it for analy- 
sis, every compound used in this part of the process must 
be separately examined for arsenic. 

Bismuth, Bi (At. Wt. 208.5). 

Occurrence, etc. — Bismuth is not abundant nor widely 
distributed in nature. It occurs for the most part native 
in veins of granite and clay slate. Among the compounds 
.of bismuth found in nature are the oxide Bi 2 3 and the 
corresponding sulphide Bi 2 S 3 . 

The ores are roasted and then treated with appropriate 
reducing agents. In different places different methods of 
extraction are employed. As the chief applications of 
bismuth are for pharmaceutical purposes, it is necessary 
that the element should be specially pure; above all, that 
it should not be contaminated with arsenic. In order to 
remove the last traces of this element the powdered bis- 
muth is generally melted with saltpetre. 

Bismuth is a hard, brittle, reddish-white substance with 
a metallic lustre. It looks very much like antimony, but 
is distinguished from it by its reddish tint. At ordinary 
temperatures it remains unchanged in the air. When 



3IO COLLEGE CHEMISTRY. 

heated to reel heat it burns with a bluish flame, forming 
the yellow oxide Bi 2 3 . 

Hydrochloric acid scarcely acts upon it; concentrated 
sulphuric acid forms bismuth sulphate, Bi 2 (SOJ 3 , in which 
the bismuth evidently plays the part of a base-forming 
element; nitric acid gives bismuth nitrate, Bi(N0 3 ) 3 , 
which is partly decomposed by water, forming so-called 
basic nitrates which are difficultly soluble in water. These 
salts will be taken up in the next chapter. 

Some bismuth is used in the preparation of alloys which 
are easily fusible, as, for example, Newton's metal, which 
contains bismuth, lead, and tin; Kose's metal, which con- 
sists of the same constituents in slightly different propor- 
tions; and Wood's metal, which consists of bismuth, lead, 
tin, and cadmium. 

Bismuth does not combine with hydrogen. 

Compounds of the Members of the Phosphorus Group 
with the Members of the Chlorine Group. — In the intro- 
duction to this chapter it was stated that the elements of 
the phosphorus group combine with chlorine in two pro- 
portions, forming compounds of the general formulas 
MC1 3 and M01 5 . Arsenic, however, forms only one com- 
pound with chlorine, AsCl 3 , while bismuth forms one of 
the formula BiCl 3 , and another, Bi 2 Cl 4 . The compounds 
of phosphorus and chlorine are the best known, and a brief 
study of these will give a fair idea of the methods of prep- 
aration and the conduct of the analogous compounds of 
the other members of the group. 

Phosphorus Trichloride, PC1 3 , is made by conducting 
dry chlorine gas upon phosphorus in a retort connected 
with a receiver. Action takes place at once with evolution 
of heat, and the trichloride distils over and is condensed 
as a liquid into the receiver. Or it can be made somewhat 
more conveniently by dissolving ordinary phosphorus in a 
little phosphorus trichloride, cooling the vessel by placing 
it in cold water, and passing dry chlorine into the solution 
until the increase in weight shows that all the phosphorus 
in solution has been transformed into the trichloride. It 



COMPOUNDS OF PHOSPHORUS AND CHLORINE. 3 1 I 

is purified by distillation on a water-bath. It is a clear, 
colorless liquid, that boils at 74°. In contact with air it 
fumes in consequence of the action of the water- vapor. 
It has a disagreeable odor of its own mixed with that ol 
hydrochloric acid. Its most characteristic decomposition 
is that which it undergoes with water, resulting in the 
formation of phosphorous and hydrochloric acids. 

The trichloride shows a strong tendency to take up 
chlorine, bromine, iodine, oxygen, and sulphur, and thus 
to become saturated as a quinquivalent element, With 
chlorine it forms the pentachloride, PCI. , w T ith oxygen the 
oxychloride, POCl 3 , and with sulphur the sulphochloride, 
PSC1 S . 

Phosphorus Pentachloride, PCL , is formed by treating 
phosphorus or the trichloride with dry chlorine. It is 
best prepared by passing chlorine through a w T ide tube 
upon the surface of the trichloride, contained in a vessel, 
which is kept cool. It is a white solid, but it generally 
has a slightly yellowish or greenish color in consequence 
of a slight decomposition into the trichloride and free 
chlorine. It sublimes below 100° without melting. When 
heated to boiling it undergoes partial decomposition into 
chlorine and the trichloride, and this decomposition is 
complete at about 300°. As the temperature is raised from 
the apparent boiling-point to the point at which the 
decomposition is complete, the color of the vapor is seen 
to grow darker in consequence of the increased quantity 
of free chlorine present. The decomposition is gradual, 
and, for any given temperature, the amount of decomposi- 
tion is constant. This kind of decomposition, which is 
known as dissociation, has been studied very carefully, and 
is found to be capable of explanation by the aid of the 
kinetic theory of gases. In a later chapter this subject 
will be treated, and a number of other examples will be 
given. Owing to this decomposition under the influence 
of heat the specific gravity of the vapor of phosphorus 
pentachloride is not what it should be, if the formula is 
PC1 5 . On the other hand, the specific gravity of the vapor 



312 COLLEGE CHEMISTRY. 

of the trichloride leads to the formula P01 3 , and that of 
the oxychloride to the formula POCl 3 . The apparent 
anomaly presented by the pentachloride is easily under- 
stood. When a molecule of the compound is converted 
into vapor, or is heated to a sufficiently high temperature, 
it is broken down in accordance with this equation : 

PC1 S = PC1 S + 01 2 . 

From the one molecule, therefore, two gaseous molecules 
are obtained. Consequently the vapor formed occupies 
twice as much space as it would if there were no decom- 
position. It follows that the specific gravity of the vapor 
must be only half what it would be if there were no 
decomposition. When the compound is converted into 
vapor in an atmosphere of phosphorus trichloride, the 
decomposition referred to does not take place, and, under 
these circumstances, the specific gravity is found to be in 
accordance with Avogadro's law, and with the formula 
PC1 5 . This case is a particularly interesting one, as it has 
played an important part in the discussions in regard to 
the validity of Avogadro's law. 

Treated with water phosphorus pentachloride first gives 
the oxychloride and hydrochloric acid, thus : 

PC1 5 + H 2 = POCl 3 + 2HC1; 

and by further action the oxychloride is converted into 
phosphoric acid, thus : 

POCI3 + 3H 2 = PO(OH) 3 + 3HC1. 

Arsenic Trichloride, AsCl 3 , is the only compound which 
arsenic forms with chlorine. It is easily made by passing 
chlorine into a retort containing powdered arsenic. 

Compounds of Antimony and Chlorine. — Antimony 
forms two compounds with chlorine, analogous in com- 
position to the two chlorides of phosphorus. These are 
antimony trichloride, SbCl^, and antimony 'pentachloride, 
SbCl 5 . The trichloride is formed by direct treatment of 



BISMUTH AND CHLORINE. 3 X 3 

the element with chlorine. It is also formed by dissolving 
antimony in hydrochloric acid with addition of nitric acid, 
and when this solution is distilled the chloride of antimony 
passes over. At the ordinary temperatures it is a solid, 
colorless, crystalline, soft substance, which, on account of 
its consistency, has received the common name " butter of 
antimony" (Butyrum Antimonii). 

Antimony chloride has a caustic action and is used in 
medicine. It is also used for the purpose of burnishing 
ironware, as gun-barrels. When treated with the chloride 
they acquire a brownish, bronze-like color. 

Pentachloride of Antimony is in general like the penta- 
chloride of phosphorus. It gives up its chlorine, however, 
somewhat more readily. 

Bismuth and Chlorine. — Bismuth differs from the other 
members of the phosphorus group in its conduct towards 
chlorine. It forms the chloride, Bi 2 Cl 4 , of which there is 
no analogue among the compounds of the other elements 
of the group with chlorine. The compound may, however, 
be regarded as analogous to the hydrogen compounds 
hydrazine, N 2 H 4 , and liquid phosphine, P 2 H 4 . Bismuth 
dichloride, Bi 2 Cl 4 , is formed by reduction of the tri- 
chloride, BiCl 3 , when the latter is treated with hydrogen 
at a temperature of about 300°. Bismuth trichloride, 
BiCl 3 , is formed by treating bismuth with chlorine, and by 
dissolving bismuth oxide, Bi 2 3 , in hydrochloric acid. 
From this solution a crystallized compound of the formula 
BiCl 3 -|- H 2 is deposited, and it is impossible to drive all 
the water off from this compound without causing decom- 
position. Treated with water the chloride is decomposed, 
forming the oxychloride, BiOCl: 

BiCl 3 + H 2 = BiOCl + 2HC1. 



314 



COLLEGE CHEMIS 77? Y. 



EXPERIMENTS. 



Phosphorus. 



Experiment 147.— [This, as well as the other experiments with 
phosphorus, should be performed only by an experienced person.] 
Arrange an apparatus as shown in Fig. 57. The neck of the 
retort is somewhat drawn out and 
bent downward and fitted air-tight 
by means of a cork to the wide glass 
tube B. Some small pieces of ordi- 
nary phosphorus are now carefully 
slipped into the retort— as much as 
is obtained by cutting up two sticks 
three to four inches long. The ap- 
paratus is then adjusted as shown in 
the figure, so that the end of the 
tube B dips below the surface of the 
water in the beaker C. The whole 
is then allowed to stand for some 
hours. The oxygen is absorbed from 
the air contained in the vessel, and 
the water rises in B. Without un- 
covering the end of B, replace the 
water in C by some that has a temperature of about 50°. Now 
heat the retort gradually, when the phosphorus will distil over 
and condense in C in the molten condition. By lowering the 
heat gradually at the end of the operation . it can finally be 
stopped without danger of breaking. 

Experiment 148. — Dissolve a little ordinary phosphorus in 
carbon disulphide. Pour some of this solution upon a strip of 
filter-paper, and let this hang in the air or wave it gently in the 
air. After the carbon disulphide has evaporated the phosphorus 
will take fire. 

Experiment 149. — Bring together in a porcelain crucible or 
evaporating-dish a little phosphorus and iodine. It will be seen 
that simple contact is sufficient to cause the two substances to act 
upon each other. Direct combination takes place, and the 
action is accompanied by light and heat. 




Fig. 



EXPERIMENTS WITH PHOSPHORUS. 



315 



Phosphorus abstracts Oxygen from other Sub- 
stances. 

Experiment 150.— Add a little of a solution of phosphorus in 
carbon disulphide to a solution of copper sulphate. What change 
takes place ? 

Experiment 151.— Put a few small pieces of ordinary phos- 
phorus in a glass tube and seal it. Heat gradually to 300°. 
Open the tube and examine the product. See whether it takes 
Are as readily as ordinary phosphorus does ; whether it dis- 
solves in carbon disulphide ; whether it melts easily when put in 
water heated to between 45° and 50°. 

Phosphine. 

Experiment 152.— Arrange an apparatus as shown in Fig. 58. 
In the small flask B put about 5 grams caustic potash dissolved 
in 10-15 cc. water, and when the solution is cold add a few small 
pieces of phosphorus the size of a pea. Pass hydrogen for some 




Fig. 58. 

time through the apparatus from the generating-flask A until all 
the air is displaced ; then disconnect at D, leaving the rubber 
tube, closed by the pinch-cock, on the tube which enters the 
flask. Gently heat the contents of the retort, when gradually a 
gas will be evolved, and will escape through the water in C. As 
each bubble comes in contact with the air it takes fire, and the 
products of combustion arrange themselves in rings, which be- 



316 



COLLEGE CHEMISTRY. 



come larger as they rise. They are extremely beautiful, particu- 
larly if the air of the room is quiet. Both the phosphorus and 
the hydrogen combine with oxygen in the act of burning. Col- 
lect some of the gas in a tube over water, and then place the 
tube mouth upward. What difference is there between the 
burning of the gas under these circumstances, and that noticed 
when the rings are formed ? Collect another tubef ul of the gas, 
and let this stand for some time. Then open the vessel by taking 
it out of the water. Has any change taken place in the gas ? 

Arsenic. 

Experiment 153. — Heat a small piece of arsenic on charcoal in 
the flame of the blowpipe. 

Arsine. 

Experiment 154. — Arrange an apparatus as shown in Fig. 59. 
Put some pure granulated zinc in the flask and pour dilute sul- 
phuric acid on it. The calcium-chloride tube serves to dry the 




Fjg. 59. 



gas. When the air is all out of the vessel and the hydrogen is 
lighted, add slowly a little of a solution of arsenic trioxide, 
As 2 3 , in dilute hydrochloric acid. The appearance of the flame 
will soon change. It will become paler, with a slightly bluish 
tint, and give off white fumes. (See next experiment.) 



EXPERIMENTS WITH ARSENIC, ETC. 3*7 



Marsh's Test for Arsenic. 

Experiment 155. — Into the flame of the burning hydrogen 
and arsine produced in the last experiment introduce a piece of 
porcelain, as the bottom of a small porcelain dish or a crucible, 
and notice the appearance of the spots. Heat by means of a 
Bunsen burner the tube through which the gas is passing, which 
should be of hard glass. Just beyond the heated place there will 
be deposited a thin layer of metallic arsenic, commonly called a 
mirror of arsenic. This deposit is due to the direct decomposi- 
tion of the arsine into arsenic and hydrogen by heat. [Compare 
ammonia, phosphine, and arsine with reference to their stability.] 

Antimony. 

Experiment 156.— Heat a small piece of antimony on char- 
coal in the blowpipe flame. Try the action of dilute and of concen- 
trated hydrochloric acid, of dilute and of concentrated nitric acid, 
and of a mixture of the two acids on a small piece of antimony. 

Stibine. 

Experiment 157. — Stibine is made by the same method as that 
used in making arsenic. Make some, using a solution of tartar 
emetic. Introduce a piece of porcelain in the flame, and after- 
wards heat the tube through which the gas is passing. Compare 
the antimony spots with the arsenic spots. Color ? Volatility ? 
Conduct towards a solution of soduim hypochlorite or hypo- 
bromite ? 

Bismuth. 

Experiment 158.— Heat a piece of bismuth on charcoal in the 
blowpipe flame. See how it conducts itself towards hydrochloric 
acid ; towards nitric acid. If a solution is obtained in either 
case, add water to it. Explain what takes place. 

Phosphorus Trichloride. 

The experiments with the chlorides of phosphorus must be 
carried on under a hood or out-of-doors. 

Experiment 159.— Arrange an apparatus as shown in Fig. 60. 
The tube A is arranged so that it can be raised or lowered in the 
retort. Put 50 to 100 grams ordinary phosphorus in the retort, 
taking precautions to prevent it from taking fire during the 



3i8 



COLLEGE CHEMISTRY. 



operation. This is best accomplished by fitting corks in both 
openings of the retort ; placing the retort in a vessel of cold 
water ; removing the cork from B ; throwing in a piece of phos- 
phorus, and quickly putting the cork in. The pieces must not be 
put in in too rapid succession. After all the phosphorus is in the 
retort, adjust the apparatus as represented, placing the receiver 
D in a dish of cold water. Now connect by means of the rubber 




Fig. 60. 



tube _2Fwith an apparatus furnishing chlorine, dried by means of 
concentrated sulphuric acid and calcium chloride. As soon as 
the chlorine comes in contact with the phosphorus action begins, 
and the product, which is phosphorus trichloride, distils over into 
the receiver. If the action is taking place too rapidly, the inside 
of the retort will become covered with a coating of red phos- 
phorus. In this case raise the tube A a little and the red coating 
will gradually disappear. If the tube is raised too high, not 
enough heat is generated, and the trichloride in the retort is con- 
verted into the pentachloride, which 
is deposited as a white coating. By 
raising and lowering the tube accord- 
ing to the indications, the retort can 
be kept clear, and all the phosphorus 
converted into the trichloride. This 
manipulation of the tube is much facili- 
tated by fitting into the cork a some- 
what larger tube, through which the 
smaller one can pass easily ; letting 
this project about an inch and a half 
above the cork and passing over it a 
piece of rubber tubing of such size 

that while the smaller tube moves through it readily, the two 
form a gas-tight joint. This is shown in Fig. 61. After the 
operation is finished, pour the liquid from the receiver into a 




EXPERIMENTS MTH PHOSPHORUS CHLORIDE. 3 X 9 

clean dry flask, and distil on a water bath. Try the action of a 
hltle of the compound on water. 



Experiment 




Fig. 62. 

ride on water, 
chloride. 



Phosphorus Pentachloride. 

160. — Put the trichloride of phosphorus obtained 
in the last experiment in a wide-mouthed bottle 
surrounded by cold water. Through a wide glass 
tube pass dry chlorine upon the surface of the 
liquid, and as the action advances, and a solid 
begins to make its appearance, stir the contents 
of the bottle. Continue the passage of the chlo- 
rine until the product is a perfectly dry solid. 
The arrangement of the bottle containing the 
trichloride, and that of the delivery-tube, is 
shown in Fig. 62. The bottle is put in a larger 
vessel containing cold water, which is renewed 
from time to time during the process. 

Try the action of a little phosphorus pentachlo- 
In a large dry flask heat a little of the penta- 



CHAPTER XVIII. 

COMPOUNDS OF THE ELEMENTS OF THE PHOS 
PHORUS GROUP WITH OXYGEN AND WITH OXYGEN 
AND HYDROGEN. 

Introduction. — The product of the direct action of 
oxygen upon phosphorus is the pentoxide P 2 5 . Arsenic, 
antimony, and bismuth, however, form the trioxides 
As 2 3 , Sb 2 3 , and Bi 2 3 . It is possible to obtain a com- 
pound of arsenic and oxygen of the formula As 2 5 , one of 
antimony, Sb 2 4 , and another, Sb 2 5 , and, finally, two 
oxides of bismuth, Bi 2 2 and Bi 2 5 . Phosphorus, further, 
forms the oxides P 4 0, P 2 3 , and P 2 4 . The table below 
contains the formulas of the above-mentioned compounds 
systematically arranged : 

P 4 Bi 2 2 

P 2 3 As 2 3 Sb 2 3 Bi 2 3 

P 2 4 Sb 2 4 

P 2 5 As 2 5 Sb 2 5 Bi 2 5 

The final products of the oxidation of the elements of this 
group, if water is present, are phosphoric, arsenic, anti- 
monic,an4 bismuthic acids. All of these are well-marked 
acMs except the last. They can all be regarded as derived 
from the normal acids of the general formula M(OH) 5 by 
loss of one or two molecules of water. The common forms 
of phosphoric, arsenic, and antimonic acids are those which 
are formed from the normal acids by loss of one molecule 
of water: 

P(OH) 5 P )(OH) 3 + H *°; 

Normal phosphoric acid Orthophosphoric acid 

320 



COMPOUNDS OF THE PHOSPHORUS GROUP. 321 



As(OH) 5 = 

Normal arsenic acid 

Sb(OH) 5 = 

Normal antimonic acid 



As 



(0 



\ (OH), 

Orthoarsenic acid 

(0 



Sb 



1 (OH), 



+ H 2 0; 
+ H 2 0. 



Orthoautimonic acid 



Bismuthic acid appears, however, to be formed from the 
normal acid by loss of two molecules of water, just as the 
so-called metaphosphoric, metarsenic, and metantimonic 
acids are : 



Bi(OH) 5 = 

Normal bismuthic acid 


Bismuthic acid 


+ 


2H 2 0; 


P(OH) 5 


P j0 2 
r (OH 

Metaphosphoric acid 


+ 


2H,0; 


As(OH) 5 = 


As |0H 

Metarsenic acid 


+ 


2H 2 0; 


Sb(OH), = 


Sb { ok 

Metantimonic acid 


+ 


2H,0. 



From the ordinary or ortho acids and from the meta 
acids more complex forms can be derived by loss of differ- 
ent quantities of water. The most common form besides 
those mentioned is that seen in the so-called pyro acids, of 
which pyrophosphoric acid is the best known example. 
It is formed from the ortho acid by loss of one molecule of 
water from two molecules of the acid, just as pyrosulphuric 
or disulphuric acid is formed from two molecules of ordi- 
nary sulphuric acid by loss of one molecule of water. The 
formation of pyrophosphoric acid from orthophosphoric 
acid takes place according to the equation 





P^ 



OH 
OH 
OH 
OH 
OH 
OH 






OH 

OH 



OH 

OH 





+ H 3 0; 



322 COLLEGE CHEMISTRY. 

or 

2H 3 P0 4 = H 4 P 2 7 + H 2 0. 

Orthophosphoric acid Pyrophosphoric acid 

Pyroarsenic and pyroantimonic acids bear the same 
relations to the ortho acids that pyrophosphoric acid bears 
to orthophosphoric acid. 

By partial oxidation of phosphorus in presence of water, 
phosphorous acid, H 3 P0 3 , is formed. The same acid is 
formed by the action of phosphorus trichloride on water. 
According to the latter method of formation we should 
expect to find that this acid is normal phosphorous acid, 
P(OH) 3 . It appears probable, however, that the acid has 

the constitution = P^OH. The acids of arsenic and 

X OH 
antimony of similar composition seem to be the normal 
acids As(OH) 3 and Sb(OH) 3 . The hydroxyl derivative of 
bismuth corresponding to these acids has no acid proper- 
ties, but on the contrary is basic. Hypophosphorous acid 
has the composition H 3 P0 2 . It is monobasic, and it 
appears therefore that it contains but one hydroxyl, as 
represented in the formula H 2 OP(OH). It is probable 
that the relations between phosphoric, phosphorous, and 
hypophosphorous acids should be represented by the 
formulas 



(OH 


f H 


( H 


OP \ OH, 


OP^ OH, 


OP^H . 


(.OH 


(oh 


(OH 


Phosphoric acid 


Phosphorous acid 


Hypophosphorous acid 



The fundamental compound, then, from which these may 
be regarded as derived is the unknown oxyphosphine 
OPH 3 . By oxidation we should expect phosphine to yield 
in successive stages the three products above named: 



H 


( H 


(H (H (OH 


H, 


OP^H, 


OP-^ H , OP^OH, OP^ OH. 


H 


(H 


{ OH (OH (OH 




Unknown 


Hypophosphorous Phosphorous Phosphoric 
acid acid acid 






PHOSPHORIC ACID. 3 2 3 

The oxidation of hydrogen sulphide takes place similarly, 
as has been shown: 

s {h> °' S {h' °' S {oH' °* s {oir 

Unknown Sulphurous acid Sulphuric acid 

With oxygen and chlorine the elements of the phos- 
phorus group form a number of compounds known as 
oxychlorides. Towards chlorine as well as towards oxygen 
all these elements except bismuth are quinquivalent. A 
part or all of the oxygen of the oxygen compounds can be 
replaced by chlorine. Starting with the chlorine com- 
pound on the one hand, oxychlorides can be obtained from 
it, until oxygen is substituted for all the chlorine, and the 
limit is reached in the oxide. So, also, hydroxyl can be 
substituted for the chlorine and the acids thus obtained. 

With sulphur phosphorus apparently forms a large 
number of compounds. Among them are two which have 
the formulas P 2 S 3 and P 2 S 5 . These plainly are analogous 
to the two oxides of phosphorus, P 2 3 and P 2 5 . 

Arsenic forms with sulphur several compounds, the 
principal of which are the disulphide, As 2 S 2 , the trisul- 
phide, As 2 S 3 , and the pe?itasnlphide, As 2 S 5 . The principal 
sulphides of antimony are those of the formulas Sb 2 S 3 and 
Sb 2 S 5 , and of bismuth those of the formulas Bi 2 S 2 and 
Bi&. 

Phosphoric Acid, Orthophosphoric Acid, H 3 P0 4 . — The 
compound to which the name phosphoric acid is generally 
applied, and from which the best known phosphates are 
derived, is that which has the formula H Q P(T. To distin- 
guish it from the other varieties it is called orthophosphoric 
acid. As has been stated, this is the final product of 
oxidation of phosphorus in the presence of water. Thus, 
when phosphorus is boiled with nitric acid it is converted 
into orthophosphoric acid; and also when phosphorus is 
burned m the air, and the product dissolved m water, 
phosphoric acid is formed. In this case the first product 
of the oxidation is the pentoxide P,^ 5? also known as 



324 COLLEGE CHEMISTRY. 

phosphoric anhydride, and when this is treated with water 
it is converted into phosphoric acid : 

P 2 5 + 3H 2 = 2H.P0,. 

The occurrence of phosphoric acid in nature has already 
been referred to in connection with the occurrence of 
phosphorus, which is found in nature almost exclusively 
in the form of phosphates, principally as calcium phos- 
phate, Ca 3 (POJ 2 , in phosphorite, apatite, and the ashes 
of bones. 

In order to prepare the acid two ways suggest them- 
selves: (1) by oxidizing phosphorus with nitric acid; and 
(2) by extracting it from one of the natural phosphates, 
as phosphorite or bone-ash. The first of these methods is 
better adapted to the preparation of pure phosphoric acid, 
such as is needed for medicinal purposes; the latter is used 
where absolute purity of the product is not required. 

The preparation of phosphoric acid from a phosphate 
is not a simple matter. If the acid were volatile or insolu- 
ble there would be no difficulty in separating it. In the 
former case it would only be necessary to proceed as in 
preparing hydrochloric and nitric acids. By adding an 
acid that is not volatile except at a high temperature, 
such, for example, as sulphuric acid, and heating, the 
non-volatile acid drives out the volatile. On the other 
hand, if phosphoric acid were insoluble in water, it could 
be separated by adding a soluble acid to one of its soluble 
salts. When, for example, nitric acid is added to a solu- 
tion of potassium tellurite, K 2 Te0 3 , tellurious acid, being 
insoluble, is thrown down: 

K 2 Te0 3 + 2HNO, = 2KN0 3 +.H 2 Te0 3 . 

But phosphoric acid is not volatile and is soluble, so that 
plainly neither of these methods can be used. By treating 
the calcium salt with sulphuric acid the calcium can be 
completely separated in the form of calcium sulphate, 
which is difficultly soluble in water and insoluble in 



PHOSPHORIC ACID. 3 2 5 

alcohol. The ideal reaction to be accomplished is that 
represented in the following equation : 

Ca 3 ( p 4 ) 2 + 3H 2 S0 4 = 3CaS0 4 + 2H 3 P0 4 . 

But when sulphuric acid is added to calcium phosphate, 
only a part of the calcium is thrown down as sulphate, the 
rest remaining in the form of primary calcium phosphate : 

Ca 3 (POJ 2 + 2H 2 S0 4 - 2CaS0 4 + CaH 4 (POJ 2 . 

The phosphate thus formed is soluble in water, and the 
calcium is not easily precipitated from it. By evaporation 
and addition of sufficient sulphuric acid and alcohol the 
precipitation can be effected, and a solution of phosphoric 
acid thus obtained. This acid is not pure, as there are 
substances in bone-ash which are not removed by the 
method described. 

Phosphoric acid for medicinal purposes is often made 
by dissolving the pentoxide in water. 

Properties. — When evaporated to the proper consistency 
the acid forms a thick syrup which slowly solidifies in the 
form of large crystals. The crystals are deliquescent. 
When heated to a sufficiently high temperature the acid 
loses water, as already explained, and yields, first, pyro- 
phosphonc, and then metaphosphoric acid. It is a tribasic 
acid, capable of yielding three classes of salts of the 
(OH (OH ( OM 

general formulas 4 OP \ OH, OF{ OM, and 0?\ OM, 
( OM ( OM ( OM 

which are known respectively as the primary, secondary, 
and tertiary phosphates. The primary and secondary 
phosphates are also known as acid phosphates, and the 
tertiary salts as neutral or normal phosphates. In these 
salts it is not necessary that all the hydrogen should be 
replaced by the same metal. There are salts in which two 
or three metals take the place of the hydrogen atoms. A 
phosphate much used in the laboratory, for example, is 
one in which one hydrogen atom of phosphoric acid is 



326 COLLEGE CHEMISTRY. 

replaced by a sodium atom, and another by the ammonium 

C OH 
group, NH . This salt has the formula OP ■! ONa , and 

( ONH 4 
is called ammonium sodium phosphate. Another phosphate 
commonly met with is ammonium magnesium phosphate, 

(ONH 4 
OP ■< ■»«- , which is derived from the acid by replace- 
ment of two hydrogen atoms in the molecule by one 
bivalent magnesium atom, and one by the ammonium 
group. 

The changes which these three classes of phosphates 
undergo when heated are of special interest. The tertiary 
phosphates are stable. The primary and secondary phos- 
phates give up all their hydrogen, which passes off in 
the form of water. Thus, primary sodium phosphate, 
H 2 NaP0 4 , loses one molecule of water from each molecule 
of the salt, and is converted into the metaphosphate, 
NaP0 3 : 

(OH 
OP \ OH = 2 P(ONa) + H 2 0. 
(ONa 

In general, the primary phosphates are converted into meta- 
phosphates by heat. 

When a secondary phosphate is heated the product is a 
pyrophosphate, as when secondary sodium phosphate is 
heated to a sufficiently high temperature it is converted 
into sodium pyrophosphate: 

( ONa , ^. AT 

0P gg. OP°g 

I OM ( , TTA 

( OH ~" i u + 2 U ? 

OP oia OP ONa 
ONa ' 0I,a 



or 



2Na 2 HP0 4 = Na 4 P 2 7 + H 2 0. 



PYROPHOSPHORIC AND METAPHOSPHORIC ACIDS. 3 2 7 

Iii general, a secondary phosphate is coverted into a 
pyrophosphate by heat. 

Pyrophosphoric Acid, H 4 P 2 7 . — When phosphoric acid 
is heated to 200°-300° until a specimen neutralized with 
ammonia gives a pure white precipitate with silver nitrate, 
it is completely transformed into pyrophosphoric acid by 
loss of water. The white precipitate referred to is the 
silver salt of pyrophosphoric acid. The silver salt of 
orthophosphoric acid is yellow. This difference in color 
led, many years ago, to a careful investigation of the 
change in composition which phosphoric acid undergoes 
when heated, and to the recognition of the existence of 
pyrophosphoric acid as distinct from orthophosphoric acid; 
and the study of the relations existing between these acids 
and metaphosphoric acid has had a strong influence in 
shaping the views of chemists in regard to the relations 
between other similar acids. The views at present held 
in regard to the relations between the common forms of 
oxygen acids and the so-called normal acids or maximum 
hydroxides are simply an extension of the ideas first intro- 
duced into chemistry by Graham in connection with the 
three varieties of phosphoric acid. The modifications of 
sulphuric acid seen in the normal acid, S(OH) 6 , the ordi- 
nary acid, S0 2 (OH) 2 , and the pyro-acid, H 2 S 2 7 , are 
examples of the same kind of relations. 

Pyrophosphates are formed, as we have seen, when the 
secondary phosphates, like disodium phosphate, HNa 2 P0 4 , 
are heated. 

Metaphosphoric Acid, HP0 3 . — This acid is formed by 
dissolving phosphorus pentoxide, P 2 5 , in cold water: 

P,0 5 + H 2 = 2HP0 3 . 

It is also formed by heating phosphoric acid to 400° : 

H.PO, = HPO, + H,0. 

Further, the metaphosphates are formed by heating the 
primary phosphates like primary sodium phosphate, 



328 COLLEGE CHEMISTRY. 

H 2 NaP0 4 . The acid is a vitreous translucent mass, known 
in the market as glacial phosphoric acid (Acidum phos- 
phoricum glacials). It is the more common commercial 
form of phosphoric acid. It is a monobasic acid, and in 
composition is analogous to nitric and chloric acids : 

HP0 3 Metaphosphoric acid. 

HN0 3 Nitric acid - 

HCIO3 Chloric acid. 

When boiled with water in which there is a little nitric 
acid metaphosphoric acid is readily converted into ortho 
phosphoric acid : 

HP0 3 + H 2 = H 8 P0 4 . 

This transformation is effected also by simply allowing the 
solution of the meta-acid in water to stand for a time, and 
by boiling the solution. 

Phosphorous Acid, H 3 P0 3 . — This acid is formed when 
phosphorus trichloride is treated with water. It is also 
formed together with phosphoric and hypophosphoric 
acids when phosphorus is allowed to lie in contact with 
moist air. The acid can be obtained from its solutions by 
evaporation, when it is deposited in transparent crystals. 

Phosphorous acid is only dibasic, its salts having the 
general formula HM 2 P0 3 . This fact has led to the belief 
that in the acid two of the hydrogen atoms are in com- 
bination with oxygen in the form of hydroxyl, while the 
third is in combination with phosphorus as represented in 

the formula OP \ OH. This conclusion finds further sup- 

(OH 
port in the conduct of some derivatives of phosphorous 
acid. 

Hypophosphoric Acid, H 4 P 2 6 , is formed together with 
phosphoric and phosphorous acids when sticks of ordinary 
phosphorus placed in glass tubes drawn out to a small 
opening at one end are exposed to the action of moist air. 



PHOSPHORIC ANHYDRIDE— PHOSPHORUS ANHYDRIDE. 3 2 9 

Hypophosphorous Acid, H 3 P0 2 , has already been re- 
ferred to, as its potassium salt is formed in the preparation 
of phosphine by the action of phosphorus upon a solution 
of potassium hydroxide : 

3KOH + 4P + 3H 2 = 3K 2 HP0 2 + PH S . 

The acid is monobasic, and this has led to the belief 
that only one of the hydrogen atoms in the molecule of the 
acid is in combination with oxygen as hydroxyl, and that 
the two others are in combination with phosphorus as 

( H 

represented in the formula OP \ H. The relation between 

(OH 
this acid and phosphorous and phosphoric acids has already 
been commented upon (see page 323). 

Phosphorus Pentoxide, Phosphoric Anhydride, P 2 5 . — 
This highest oxidation-product of phosphorus is formed 
by burning the element in air or in oxygen. It is a white 
powder that attracts moisture from the air and becomes 
liquid, This power to combine with water is its most 
characteristic property. It forms first, as we have seen, 
metaphosphoric acid and, by further action, orthophos- 
phoric acid. Its action towards water is strongly sugges- 
tive of the action of sulphur trioxide or sulphuric 
anhydride towards water. Owing to this power to combine 
with water, phosphorus pentoxide is used for the purpose 
of drying gases and as a dehydrating agent. 

Phosphorus Trioxide, or Phosphorus Anhydride, P 2 3 
(or P^Og), is formed by burning phosphorus so that the 
air does not have free access to it, as by putting a piece of 
phosphorus in a glass tube drawn out to a fine opening, 
drawing air over the phosphorus, and warming it gently. 
In this way not enough air can get access to the phos- 
phorus to convert it into the pentoxide. The trioxide has 
such a strong tendency to pass over into the pentoxide 
that when brought into the air it takes fire and burns, 
forming the higher oxide. It is readily converted into 
phosphorous acid by water. 



33° COLLEGE CHEMISTRY. 

Phosphorus Oxychloride, P0C1 3 . — This compound has 
been referred to in connection with the chlorides of phos- 
phorus, It is formed by the action of water upon the 
pentachloride : 

PC1 5 + H 2 = POCI3 + 2HC1. 

It is a liquid that boils at 107.2°. 

Arsenic Acid, H 3 As0 4 . — The compound of arsenic and 
oxygen which is most readily obtained is the trioxide, 
As 2 3 , and this is formed by direct combination of the 
two elements. When this is oxidized either with aqua 
regia or by passing chlorine into water in which the tri- 
oxide is suspended it is converted into arsenic acid : 

As 2 3 + 3H 2 + 20 = 2H 3 As0 4 . 

From its solutions it is obtained in crystallized form. 
According to the temperature to which it is heated the 
deposit has the composition of the ortho-acid, H 3 As0 4 , of 
the pyro-acid, H 4 As 2 7 , or of the meta-acid, HAs0 3 . 
Perfect analogy with the phosphorus compounds is here 
observed. When the pyro- and meta-acids are dissolved in 
water they are converted into the ortho-acid. Arsenic 
acid, like phosphoric acid, is a strong tribasic acid, form- 
ing three series of salts which under the influence of heat 
conduct themselves like the corresponding phosphates, the 
primary salts yielding pyro-arsenates, and the secondary 
salts yielding meta-arsenates. 

Arsenious Acid, H 3 As0 3 , is not known in the free state, 
but salts derived from it are formed by treating arsenic tri- 
oxide with bases. 

Arsenic Trioxide, As 2 3 (As 4 6 ). — This compound is 
commonly called arsenic or white arsenic. It is the most 
important of all the compounds of the element arsenic. 
It finds applications for many purposes, and is manufac- 
tured in large quantities. It occurs in small quantity in 
nature, but that which comes into the market is manufac- 



ANSEN1C TRIOXIDE. 33 * 

tnred by roasting natural arsenides, particularly arsenical 
pyrites, FeAsS. The products of roasting this compound 
are ferric oxide, Fe,,0 3 , sulphur dioxide, S0 2 , and arse lie 
trioxide, As 2 O s . Of these, the first is a non-volatile solid, 
the second a gas, and the third a volatile solid. By pass- 
ing the volatile products through properly constructed 
canals the arsenic trioxide is condensed on the walls. 
Some of the powder thus obtained must be subjected to a 
second process of distillation to make it pure enough for 
the market. In a recent year over 6000 tons of this sub- 
stance were produced in England and Saxony. 

Arsenic trioxide is a colorless, amorphous, vitreous mass. 
Gradually it becomes opaque and crystalline, with an 
appearance like that of porcelain. It crystallizes in two 
forms, the common one being that of regular octahedrons. 
Under exceptional conditions it crystallizes in the form of 
rhombic prisms. When heated it sublimes, and is 
deposited on a cold surface in the form of octahedrons. 
Arsenic trioxide is difficultly soluble in water, but more 
easily in hydrochloric acid. 

The trioxide is easily reduced. When heated with 
potassium cyanide, KON, or with charcoal in a dry glass 
tube arsenic is deposited above the flame in the form of a 
dark lustrous layer. When brought into a vessel from 
which hydrogen is being evolved it is reduced to arsine. 

The specific gravity of the vapor of the oxide shows that 
it has the formula As 4 6 , and not As 2 3 ; as, however, 
most of its reactions can be more conveniently expressed 
by the aid of the simpler formula, the latter is commonly 
used. 

Arsenic trioxide has a weak, disagreeable, sweet taste, 
and is an active poison. A dose of from two to three 
grains is sufficient to cause death unless it is ejected by 
vomiting, or rendered harmless by being converted into an 
insoluble compound. It is possible, by beginning with 
small doses, and gradually increasing them, to accustom 
the human body to considerably larger doses than that 
mentioned. It strengthens the power of the respiratory 



33 2 COLLEGE CHEMISTRY. 

organs, and consequently facilitates mountain-climbing. 
The peasants in some mountain regions are said to use it 
habitually. It is much used in medicine, especially in skin 
diseases. It is also used extensively as a rat-poison. The 
most efficient antidote is a mixture of ferric hydroxide, 
Fe(OH) 3 , and magnesia, which forms with arsenic trioxide 
an insoluble compound. 

Arsenic Pentoxide, As 2 5 , is formed by igniting arsenic 
acid. If heated too high the pentoxide breaks down into 
arsenic trioxide and oxygen. A marked difference will be 
observed between the conduct of the oxides of phosphorus 
and that of the corresponding oxides of arsenic. While 
phosphorus trioxide takes up oxygen spontaneously when 
exposed to the air, and the pentoxide is not decomposed 
by heat, the trioxide of arsenic does not under any circum- 
stances take up oxygen directly, and the pentoxide easily 
breaks down into the trioxide and oxygen when heated. 

Sulphides. — There are three compounds of arsenic with 
sulphur — the disulphide, As 2 S 2 , the trisulphide, As 2 S 3 , 
and the pentasulphide, As 2 S 5 . 

Arsenic Disulphide, As 2 S 2 , occurs in nature and is 
known as realgar. It can also be obtained by melting 
arsenic and sulphur together in the right proportions. It 
forms an orange-red powder which was formerly used as a 
pigment. 

Arsenic Trisulphide, As 2 S 3 , is found in nature and is 
called orpiment or king's yellow. It can be prepared by 
melting together arsenic and sulphur in the proper pro- 
portions, and by precipitating a solution of arsenic trioxide 
in hydrochloric acid with hydrogen sulphide. It melts, 
forming a red liquid. The natural substance, as well as 
that which is precipitated by means of hydrogen sulphide, 
is yellow. It dissolves in soluble sulphides, forming salts 
of sulpharsenious acid, H 3 AsS 3 , or HAsS 2 . The salts are, 
for the most part, derived from the acid of the latter 
formula. There is, therefore, perfect analogy between the 
oxygen and sulphur compounds, for, as we have seen, when 
arsenic trioxide is dissolved in potassium hydroxide a salt 



ANTIMONIC ACID— ANTIMONY TRIOXIDE. 333 

of the formula KAs0 2 is formed. The analogy is clearly 
shown by means of the equations 

As 2 3 + 2KOH = 2KAs0 2 + H 2 0; 

As 2 S 3 + 2KSH = 2KAsS 2 + H 2 S. 

Arsenic Pentasulphide, As 2 S 5 , is formed by melting 
sulphur and arsenic together in the proper proportions, 
and by precipitating a solution of sodium sulpharsenate 
with hydrochloric acid : 

2Na 3 AsS 4 + 6HC1 = 6NaCl + As 2 S 5 + 3H 2 S. 

Antimonic Acid, H 3 Sb0 4 . — This acid is the final product 
of the oxidation of antimony when treated with aqua regia. 
It need only be said that it is very similar to phosphoric 
and arsenic acids; and that, like these, it yields a meta- 
and a pyro-acid of the formulas HSb0 3 and H 4 Sb 2 7 . The 
acid of the formula OSb(OH) 3 , or orthoantimonic acid, is 
known in the free state, and is formed by treating a soluble 
salt of antimonic acid with sulphuric or nitric acid : 

OSb(OK) 3 + 3HN0 3 = 3KN0 3 + OSb(OH) 3 . 

An acid Sb 2 0(OH) 8 is also known in the free state, being 
formed by the action of antimony pentachloride upon 
water. The lower oxides of antimony, the trioxide, 
Sb 2 3 , and the tetroxide, Sb 2 4 , are not strongly acidic; 
that is to say, they do not readily form salts when treated 
with bases. In this respect the trioxide of antimony 
differs markedly from the corresponding oxides of phos- 
phorus and arsenic. 

Antimony Trioxide, Sb 2 3 . — This compound is found in 
nature as white ore of antimony, and is easily formed by 
burning antimony in the air and by oxidizing it with nitric 
acid or saltpetre. That formed by burning antimony in 
the air always contains some of the tetroxide, and by heat- 
ing it long enough in the air and to a temperature high 
enough it is completely transformed into the tetroxide. 



334 COLLEGE CHEMISTRY. 

When the trioxide is dissolved in caustic soda a salt of the 
formula NaSb0 2 is formed. This is plainly derived from 
an acid of the formula HSb0 2 , which bears a simple rela- 
tion to normal antimonious acid. Towards most bases, 
however, antimony trioxide does not conduct itself as an 
acid. On the other hand, towards the stronger acids it 
acts as a base. 

Salts of Antimony. — The salts of antimony are derived 
either from the hydroxide Sb(OH) 3 or from the hydroxide 
SbO.OH. The salts of the first class are called antimony 
salts ; those of the second class are called antimonyl salts. 
In the salts formed when the trihydroxide of antimony is 
completely neutralized by acids, the antimony takes the 
place of three atoms of hydrogen. Thus, the nitrate has 
the formula Sb(N0 3 ) 3 ; the sulphate has the formula 
Sb 2 (S0 4 ) 3 ; etc. Besides these normal salts there are, how- 
ever, basic salts. Thus there are two basic nitrates possi- 

( OH (OH 

ble of the formulas Sb \ OH and Sb \ N0 3 . 
(NO, (NO, 

The formation of antimonyl salts is illustrated by the 
sulphate. This may be regarded as formed by the action 
of sulphuric acid upon the hydroxide SbO.OH, which is 
analogous in composition to the acid of arsenic of the 
formula AsO. OH: 



2SbO. OH + gg > S0 2 = g£g; g > S0 2 + 2H 2 0. 



The product is antimonyl sulphate. The weak basic char- 
acter of the hydroxides of antimony is shown by the fact 
that many of its salts are decomposed by water. The salt 
of antimony which is most commonly met with is the 
so-called tartar emetic, which appears to be an antimonyl 
potassium salt of tartaric acid. Tartaric acid is a dibasic 

( OH 

acid of the formula CJ3.fi 4 < ^tt. When one of its acid 

hydrogen atoms is replaced by potassium, and the other 



ANTIMONY COMPOUNDS. 335 

by the antimonyl group SbO, the salt thus formed is tartar 
emetic, C 4 H 4 4 J , Q-g- . 

Antimony trioxide dissolves in hydrochloric acid, form- 
ing the trichloride, and this, as has been stated, is decom- 
posed by water yielding oxy chlorides. 

Antimony Tetroxide, Sb 2 4 . — This compound is most 
easily obtained by igniting antimonic acid, H 3 Sb0 4 . Two 
reactions are of course involved : 



2H 3 Sb0 4 =: Sb 2 5 + 3H 2 0; 
Sb 2 5 = Sb 2 4 + 0. 

It is also formed by igniting the trioxide in the air. At 
ordinary temperatures the tetroxide is white, but it 
becomes yellow when heated. Towards strong acids this 
oxide acts like a weak base. A potassium salt of the 
formula K 2 Sb 2 5 is known, which is derived from the acid 
H 2 Sb 2 5 , and this in turn from the simpler acid SbO(OH) 3 
by loss of water. The oxide itself is regarded by some as 
an antimonyl salt of metantimonic acid, Sb0 2 .OH, of the 
formula Sb0 2 .O.SbO. 

Antimony Pentoxide, Sb 2 5 . — The tetroxide of antimony 
does not combine with oxygen to form the pentoxide. 
The latter can be obtained on]y by gentle ignition of anti- 
monic acid, care being taken not to raise the temperature 
high enough to decompose the pentoxide into the tetroxide 
and oxygen. The fact that the pentoxide readily yields 
salts of antimonic acid when treated with basic solutions 
was mentioned under Antimonic Acid. 

Antimony Trisulphide, Sb 2 S 3 . — This compound occurs in 
nature in considerable quantity and is the chief source of 
antimony. It is known as stibniie and antimony blende. 
In some localities, especially in Japan, it occurs in large 
crystals of great beauty. When heated in the air, or 
roasted, it is converted into the trioxide, and finally into 
the tetroxide, while the sulphur escapes as the dioxide, 



33 6 COLLEGE CHEMISTRY. 

Hydrochloric acid dissolves the trisulphide in the form of 
the chloride with evolution of hydrogen sulphide : 

Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. 

Nitric acid converts it into the oxide with separation of 
sulphur. When a solution of antimony chloride is treated 
with hydrogen sulphide, the trisulphide is thrown down. 
This artificially-prepared trisulphide has an orange-red 
color, while that which occurs in nature is black or gray. 
The sulphide dissolves in solutions of metallic sulphides, 
forming salts of sulphantimonious acid, either SbS.SH or 
Sb(SH) 3 . 

Antimony Pentasulphide, Sb 2 S 5 , is formed by passing 
hydrogen sulphide into a solution of antimonic acid or by 
decomposing a salt of sulphantimonic acid by means of an 
acid. The action takes place thus : 

2H 3 Sb0 4 + 5H 2 S = Sb 2 S 5 -j- 8H 2 0; 
2Na 3 SbS 4 + 6HC1 = 6NaCl + Sb 2 S 5 + 3H 2 S. 

It is, when dry, a golden-yellow powder known as sulphur 
auratum. It dissolves easily in solutions of metallic sul- 
phides, forming the sulphantimonates, of which the sodium 
salt, Na 3 SbS 4 , known as Schlippe's salt, is a good example. 
The action is represented by this equation : 

Sb 2 S 5 + 6NaSH = 2Na 3 SbS 4 + 3H 2 S. 

When heated in the air the pentasulphide gives off 
enough sulphur to form the trisulphide; while when the 
pentoxide is heated it is converted into the tetroxide. 
The sulphantimonates are decomposed when treated with 
acids and the pentasulphide is thrown down. 

Oxychlorides of Antimony. — Under the head of Anti- 
mony Trichloride the fact was mentioned that this com- 
pound is decomposed by cold water as represented in the 
equation 

SbCl 3 + H 2 = SbOCl + 2HC1. 



OXIDES AND SALTS OF BISMUTH. 337 

If, however, hot water is used, the composition of the 
product approximates to that represented by the formula 
Sb 4 5 Cl 2 . This complex mixture of oxychlorides is known 
as the li Powder of Algaroth." It may be regarded as 
derived from the simple oxychloride by loss of antimony 
trichloride, thus: 

5SbOCl = Sb 4 5 Cl 2 + SbCl s . 

Many other oxychlorides besides the two mentioned have 
been obtained, but they are all more or less closely related 
to the simple compound SbOCl. 

Oxides of Bismuth. — The principal compound of bis- 
muth and oxygen is the trioxide, Bi 2 O s , which is formed 
when bismuth is burned in the air. It is a yellow powder. 
It is basic and forms salts which in composition correspond 
to the salts of antimony. Like the latter, they are of two 
classes — the bismuth salts and the bismuthyl salts. The 
former are derived from the triacid base, Bi(OH) 3 , the 
latter from the monacid base, BiO(OH). 

Salts of Bismuth. — The best known salts of bismuth are 
those which it forms with sulphuric and with nitric acids. 
There is a sulphate of the formula BiII(SOJ 2 formed by 
dissolving bismuth oxide in dilute sulphuric acid. The 
sulphate which is most stable in the presence of water is 
the bismuthyl salt, (BiO) 2 S0 4 . When bismuth is dissolved 
in nitric acid and the solution evaporated to dryness the 
salt Bi(N0 3 ) 3 -|- 10H 2 O is obtained. This salt is decom- 
posed when heated, and by water, forming basic nitrates 
of bismuth. The composition of the basic nitrate obtained 
by decomposing the neutral nitrate with water differs 
according to the conditions. Hot and cold water produce 
different results. A solution containing much nitric acid 
does not give the same result as one which contains little, 
etc. As basic bismuth nitrate is used in medicine it is 
necessary that specific directions should be given for its 
preparation, in order that a substance of the same com- 
position should always be obtained. The basic nitrate of 



338 COLLEGE CHEMISTRY. 

bismuth, or the subnitrate, as it is frequently called in 
pharmacy, is much used in medicine as a remedy in 
dysentery and cholera. It is also used as a cosmetic. 

Bismuth Dioxide, Bi 2 2 , is formed as a brown precipi- 
tate when potassium hydroxide is added to a solution of 
bismuth chloride and stannous chloride, SnCl 2 . 

Bismuth Pentoxide, Bi 2 5 , is formed by oxidizing the 
trioxide, by means of chlorine, in alkaline solution. 
Although some experimenters appear to have obtained salts 
of bismuthic acid, as, for example, KBi0 3 , others have 
failed to obtain them. In any case it is evident that the 
acid properties of the oxide are very weak. 

Bismuth Trisulphide, Bi 2 S 3 , occurs in nature, and is 
formed by precipitating bismuth from solutions of its salts 
with hydrogen sulphide. It dissolves in hot concentrated 
hydrochloric acid and in nitric acid. It does not dissolve 
in solutions of the sulphides as the sulphides of arsenic 
and antimony do. 

Bismuth Oxychloride, BiOCl, which in composition is 
analogous to the simplest form of antimony oxychloride, 
is thrown down as a white powder when a solution contain- 
ing bismuth chloride is treated with water: 

Bi01 3 + H 2 = BiOCl + 2H01. 

Family V, Gboup A. 

As the members of Group A, Family VII, are related 
to Group B of the same family; and as the members of 
Group A, Family VI, are related to the members of 
Group B of the same family, so the members of Group A, 
Family V, are related to the members of Group B, which 
have just been studied. The members of Group A are 
vanadium, columbium, tantalum, and didymium, all of 
which are rare. Of these vanadium has been most 
thoroughly investigated, and columbium next. 

Vanadium, V (At. Wt. 51.2). — This element occurs in 
nature in the form of vanadates or salts of vanadic acid, 
H 3 V0 4 , which is analogous to phosphoric acid. 



BORON. 339 

Vanadic Acid, H 3 V0 4 , is the most important and best 
known of the compounds of vanadium. It is the final 
product of the oxidation of vanadium, and bears to this 
element the same relation that phosphoric, arsenic, and 
antimonic acids bear to phosphorus, arsenic, and anti- 
mony. The vanadates are derived from ortho-, meta-, 
and pyrovanadic acids, though the most stable ones are 
the metavanadates, MV0 3 . The free metavanadic acid is 
known. It is a beautiful golden-yellow compound, which 
may be used as a substitute for gold bronze. 

Columbium, Cb (At. Wt. 94). — This element, which is 
sometimes called niobium, occurs in the mineral colum- 
bite. 

Tantalum, Ta (At. Wt. 183). — Tantalum occurs in the 
minerals columbite and tantalite, accompanied by niobium. 

Didymium consists of two very similar elements, neody- 
mium and praseodymium. In some of their compounds 
they show a resemblance to the members of this group. 
They form, for example, an oxide of the formula Di 2 5 . 
On the other hand, they seem to be more closely related 
to cerium and lanthanum, which are also very rare ele- 
ments, occurring associated with didymium. These will 
be further treated of in connection with lanthanum and 
cerium. 

Bolton, B (At. Wt. 11). 

General. — Although the element boron i3 not a member 
of the family to which nitrogen and phosphorus belong, 
it nevertheless resembles the members of this family in 
some respects. It belongs to the same family as aluminium, 
and in the composition of its compounds it is undoubtedly 
similar to aluminium ; but, on the other hand, its oxide is 
distinctly acidic, while that of aluminium is basic. 

Occurrence. — Boron occurs in nature chiefly in the form 
of boric acid, or as salts of this acid, particularly a sodium 
salt known as borax. 

Preparation. — From borax and the other borates the acid. 



340 COLLEGE CHEMISTRY. 

can easily be obtained. When heated, water is given off, 
and boron trioxide, B 2 3 , is left : 

2B(OH) 3 = B 2 O s + 3H 2 0. 

By heating the oxide with potassium amorphous boron is 
obtained. By melting the oxide with aluminium, boron 
is formed and is dissolved in the molten aluminium, from 
which, on cooling, it is deposited in crystals. Amorphous 
boron in almost pure form is obtained by heating borax 
with magnesium powder. One of the chief difficulties 
encountered in preparing boron is to prevent the element 
from combining with the nitrogen of the air. At the high 
temperature at which the reduction takes place the two 
elements combine very readily to form the compound 
boron nitride, BN. The crystals obtained in the process 
described are not pure boron, but contain aluminium, or 
carbon and aluminium, apparently in combination with 
the boron. The crystals are very hard, and some of them 
have a high lustre. 

Properties. — Amorphous boron is a greenish-brown 
powder. It burns when heated in the air or in oxygen, 
the product being the trioxide B 2 3 . Strong oxidizing 
agents, like nitric acid and saltpetre, readily oxidize it, 
forming boric acid. It combines readily also with many 
other elements, as with chlorine, nitrogen, and sulphur. 
When it is brought into the melting hydroxides or car- 
bonates of potassium or sodium, it forms borates o± the 
corresponding metals. 

Boron Trichloride, BC1 3 . — This compound is formed by 
heating boron in a current of dry chlorine, and by heating 
a mixture of boron trioxide and charcoal in chlorine : 

2B 2 3 + 3C + 6C1 2 = 4B01 3 + 3C0 2 . 

This reaction is especially interesting on account of its 
double character. Carbon alone could not reduce the 
boron trioxide at the temperature employed; nor could 
the chlorine alone displace the oxygen and form the 



BORIC ACID. 341 

chloride, but when both chlorine and carbon act together 
these changes take place, one aiding the other. 

The chloride is a liquid that boils at 17°. Like phos- 
phorus trichloride, it is easily decomposed by water, 
forming boric acid, which, as will be seen, is analogous in 
composition to phosphorous acid and arsenious acid: 

BCI3 + 3H 2 = B(OH) 3 + 3HC1. 

Boric Acid, B(OH) 3 . — Boric acid occurs free in nature 
and in the form of salts, of which the principal one is 
borax. Besides borax, which is a sodium salt derived from 
tetraboric acid, H 2 B 4 7 , there are other natural borates, 
as boracite, which is a magnesium salt combined with 
magnesium chloride; and datholite, which is made up of 
silicic acid, boric acid, and the element calcium. One of 
the most interesting natural forms of boric acid is that 
which is given off from the earth with steam. Such jets 
of steam are met with in many volcanic regions, and are 
called fumaroles. In Tuscany many of the fumaroles are 
charged with small quantities of boric acid, which is 
somewhat volatile with steam. Those at Monte Cerboli 
and Monte Rotundo in Tuscany are utilized for the pur- 
pose of obtaining the boric acid. For this purpose basins 
are built over the fumaroles and filled with water, so that 
the steam is condensed and the boric acid dissolved in the 
water. The solutions formed at the higher levels flow into 
basins at lower levels, and finally become charged with a 
considerable quantity of the acid, when it is evaporated to 
crystallization by the aid of the heat furnished by the 
fumaroles. The acid obtained in this way is not pure, but 
it can be purified by recrystallization. 

Boric acid can also be made from borax by heating the 
salt in solution with dilute sulphuric acid : 

Na,B 4 T + H 2 S0 4 -f 5H 2 = Na 2 S0 4 + 4B(OH) 3 . 

If the solution is sufficiently concentrated the boric acid 
crystallizes out on cooling. 

Boric acid is easily soluble in water, and crystallizes 



342 COLLEGE CHEMISTRY. 

from the solution. It is also soluble in alcohol, and this 
solution burns with a characteristic green flame. The acid 
is quite volatile with water-vapor. When heated at 100° 
orthoboric acid loses one molecule of water, and is con- 
verted into metaboric acid, HB0 2 ; at 160° it yields tetra- 
loric acid, H 2 B 4 7 ; and at a higher temperature it is 
converted into boron trioxide or boric anhydride, B 2 3 . 
These changes are represented in the equations following : 

H 3 B0 3 = HBO J + H ! 0; 
4H 3 B0 3 = H,B 4 0, + 5H 2 0. 

The most stable salts are the tetraborates and meta- 
borates. Borax is the sodium salt of tetraboric acid, 
Na 2 B 4 O r The salts of orthoboric acid are unstable. 
They break down when treated with water, forming free 
boric acid and either metaborates or tetraborates. 

Most of the boric acid obtained from Tuscany is used 
in the manufacture of borax, a salt which finds extensive 
application. 

Salts of Boron. — Although the most characteristic com- 
pounds of boron are those in which it acts as an acid- 
forming element, it forms some compounds in which its 
power as a base-former is shown. Thus, with concen- 
trated sulphuric acid the trioxide forms a compound 
which appears to be pyrosulphuric acid, H 2 S 2 7 , in which 
the group BO, analogous to antimonyl, SbO, is sub- 
stituted for one hydrogen. It has the composition 
(BO)HS 2 7 . Further, when concentrated phosphoric acid 
acts upon crystallized boric acid, boron phosphate, BP0 4 , 
is formed. This compound is characterized by great 
stability. Concentrated acids, for example, do not decom- 
pose it. It also forms a salt which appears to be analogous 
to tartar emetic, which, as has been pointed out, is prob- 
ably antimonyl potassium tartrate, C 4 H 4 6 j -g- . This is 

{BO 
g- , which 

may be called boryl potassium tartrate. 



EXPERIMENTS WITH PHOSPHORIC ACID, ETC. 343 

Nitrogen Boride, BN. — This compound has been referred 
to in connection with the preparation of boron. It is 
easily obtained by igniting a mixture of dehydrated borax 
and ammonium chloride. It forms a white powder, which 
is insoluble in water, and is characterized by great stability. 
At red heat it is decomposed by water-vapor into ammonia 
and boric acid: 

2BN + 6H 2 = 2B(OH) 3 + 2NH 3 . 



EXPERIMENTS. 
Phosphoric Acid. 
Experiment 161. — In a flask connected with an inverted con- 
denser, as shown in Fig. 63, boil 10 to 15 grams of ordinary phos- 




Fig. 63. 

phorus with 250 cc. ordinary commercial nitric acid. If neces- 
sary, add more acid after a time. Boil gently until the phos- 
phorus disappears. Evaporate the solution to complete dryness, 
so as to get rid of all the nitric acid. 

Experiment 162. — Try the action of a solution of ordinary 
sodium phosphate on a solution of silver nitrate. Heat a little of 
the phosphate in a porcelain crucible to redness. After cooling, 
try the action of the salt left in the crucible on silver nitrate. 

Arsenic Acid. 
Experiment 163. — Pass chlorine into water containing arsenic 
trioxide in suspension, until the oxide is dissolved. Evaporate to 



344 COLLEGE CHEMISTRY. 

crystallization. Into a dilute solution of the product thus ob- 
tained, to which some hydrochloric acid is added, pass hydrogen 
sulphide. Explain the changes. 

E EDUCTION" OF ARSENIC TRIOXIDE. 

Experiment 164. — In the bottom of a dry tube of hard glass of 
the form represented in Fig. 64 put a minute piece of arsenic tri- 
oxide, and just above it a small bit of charcoal. Heat 
gently. Explain the change. 

Sulphides of Arsenic. 
Experiment 165. — Pass hydrogen sulphide into a di- 
lute solution of arsenic trioxide in hydrochloric acid. — 
Filter off the precipitate, and try the action of ammo- 
nium sulphide on some of it. 

f Sulphides of Antimony. 

Experiment 166. — Pass hydrogen sulphide into a so- 
lution of antimonic acid made by treating antimony 
Fig. 64. with aqua regia and diluting with water. Pass hydro- 
gen sulphide into a solution of antimony trichloride made by 
dissolving stibnite or antimony trisulphide in hydrochloric acid. 
Try the action of ammonium sulphide on the precipitates after 
filtering. 

OXYCHLORIDES OF ANTIMONY. 

Experiment 167. — Treat a solution of antimony trichloride with 
water. 

Basic Nitrates of Bismuth. 

Experiment 168. — Dissolve a little bismuth in nitric acid and 
evaporate. Add water. 

Boron. 

Experiment 169. — Make a hot solution of 30 grams crystallized 
borax in 120 cc. water. Add slowly 10 grams concentrated sul- 
phuric acid. On cooling, the boric acid will crystallize out. 
What evidence have you that the substance which crystallizes out 
of the solution is not borax ? Try the solubility in alcohol of 
specimens of each. Is there any difference ? Treat a few crys- 
tals of borax with about 10 cc. alcohol ; pour off the alcohol and 
set fire to it. Treat a few crystals of the boric acid in the same 
way. What difference do you observe ? Distil an aqueous solu- 
tion of boric acid, and determine whether any of the acid passes 
over with the water-vapor. 



CHAPTER XIX. 

CARBON (C, At. Wt. 12) AND ITS SIMPLER COMPOUNDS 
WITH HYDROGEN AND CHLORINE. 

Introductory. — Carbon bears to Family IV relations 
similar to those which nitrogen, oxygen, and fluorine bear 
to Families V, VI, and VII. Towards hydrogen, as well 
as towards chlorine and oxygen, carbon is quadrivalent, 
and towards oxygen it is also bivalent. In this family the 
maximum oxygen-valence coincides with the hydrogen- 
valence, while, as has been seen, in Families V, VI, and 
VII, the oxygen-valence is higher than the hydrogen- 
valence, the difference becoming greater from Family V 
to VII. While the higher oxygen compounds of Family 
IV are acidic, forming acids which are derived from the 
normal acid, R(OH) 4 , the lower oxides are not generally 
acid. The hydrogen compounds of the general formula 
MH 4 , of which there are but two, those of carbon and 
silicon, have neither acid nor basic properties. Carbon is 
distinguished by the large number of the compounds into 
which it enters, all of which are more or less closely 
related to a comparatively small number of fundamental 
forms. Silicon also forms a large number of compounds, 
as we shall see; but these are of a different kind from 
those obtained from carbon. 

Occurrence of Carbon. — In general, substances which 
are obtained from the vegetable or animal kingdom blacken 
when heated to a sufficiently high temperature, and after- 
wards, if they are heated in the air, they burn up, as we 
say. When we consider the great variety of substances 
found in living things, it appears remarkable that nearly 

345 



346 COLLEGE CHEMISTRY. 

all have this property in common. It is due to the fact 
that nearly all animal and vegetable substances contain the 
element carbon. When they are heated the other ele- 
ments present are first driven off in various forms of com- 
bination, while the carbon is the last to go. Hydrogen 
and oxygen pass off as water; hydrogen and nitrogen as 
ammonia; and much of the carbon also passes off in com- 
bination with hydrogen, with hydrogen and oxygen, and 
with nitrogen and hydrogen. If the heating is carried 
on in the air, the carbon finally combines with oxygen 
to form a colorless gas — it burns up. Carbon is the cen- 
tral element of organic nature. There is not a living 
thing, from the minutest microscopic animal to the mam- 
moth, from the moss to the giant tree, which does not 
contain this element as an essential constituent. The 
number of the compounds which it forms is almost in- 
finite, and they present such peculiarities that they are 
commonly treated of under a separate head, " Organic 
Chemistry " There is no good reason for this, except the 
large number of the compounds. For our present purpose 
it will suffice to consider the chemistry of the element 
itself, and of a few of its more important simple com- 
pounds. 

From what has already been said, it will be seen that 
the principal form in which carbon occurs in nature is in 
combination with other elements. It occurs not only in 
living things, but in their fossil remains, as in coal. 
Coal-oil, or petroleum, the formation of which is perhaps 
due to the action of water on metallic carbides, consists 
of a large number of compounds which contain only 
carbon and hydrogen. Most products of plant-life contain 
the elements carbon, hydrogen, and oxygen. Among the 
more common of these products may be mentioned sugar, 
starch, and cellulose. Most products of animal life con- 
tain carbon, hydrogen, oxygen, and nitrogen. Among 
them may be mentioned albumen, fibrin, casein, etc. 
Carbon occurs in the air in the form of carbon dioxide. 
It also occurs in the form of salts of carbonic acid; the 



DIAMOhlD-GRAPHlTE. 347 

carbonates, which are very widely distributed, forming 
whole mountain ranges. Limestone, marble, and chalk 
are varieties of calcium carbonate. 

Uncombined, the element occurs pure in two very 
different forms in nature: (1) as diamond; and (2) as 
graphite, or plumbago. 

Diamond. — The diamond occurs in few places on the 
earth, and but little is known as to the conditions which 
gave rise to its formation. The celebrated diamond beds 
are in India^ Borneo, Brazil, and South Africa. When 
found, diamonds are generally covered with an opaque 
layer, which must be removed before its beautiful proper- 
ties are apparent. The crystals are sometimes regular 
octahedrons, though usually they are somewhat more 
complicated, and the faces are frequently curved. It is 
the hardest substance known. For use as a gem it must 
be cut and polished. The object in view is to bring out 
as strikingly as possible its brilliancy by exposing the faces 
favorably to the action of the light. If heated to a very 
high temperature without access of air, it swells up and 
is converted into a black mass resembling coke. The 
change takes place without loss in weight. Heated to a 
high temperature in oxygen, it burns up, yielding only 
carbon dioxide. It is insoluble in all known liquids at 
ordinary temperatures. It dissolves, however, in molten 
cast iron and in some other molten metals. Small 
diamonds have recently (1897) been made by Moissan by 
dissolving carbon in cast iron with the aid of an electric 
furnace, and suddenly cooling the mass. Under these 
conditions the carbon in the inner parts of the mass is 
under great pressure while passing into the solid form 
from solution. 

Graphite. — Graphite, or plumbago, is found in nature 
in large quantities. Sometimes it is crystallized, but in 
forms entirely different from those assumed by the 
diamond. It can be prepared artificially by dissolving 
charcoal in molten iron, from which solution, on cooling, 
graphite is deposited. It can also be prepared by heating 



34^ COLLEGE CHEMISTRY. 

charcoal to a high temperature in the electric furnace. 
This variety differs in some of its properties from that 
which crystallizes from iron. Both varieties are found in 
nature. It has a grayish-black color and a metallic lustre. 
It is quite soft, leaving a leaden-gray mark on paper when 
drawn across it, and it is hence used in the manufacture 
of so-called lead -pencils. It is sometimes called black-lead. 
When heated without access of air it remains unchanged. 
Heated to a very high temperature in the air, or in oxygen, 
it burns up, forming only carbon dioxide. Like the 
diamond, it is insoluble in all known liquids at ordinary 
temperatures. 

Amorphous Carbon. — All forms of carbon which are 
not diamond, nor graphite, are included under the name 
amorphous carbon. The name signifies simply that it is 
not crystallized. The most common form of amorphous 
carbon is ordinary charcoal. 

Charcoal is that form of carbon which is made by the 
charring process. This consists simply in heating wood 
without a sufficient supply of air to effect complete com- 
bustion. The substance almost exclusively used in the 
manufacture of charcoal is wood. As has already been 
stated, wood is made up of a large number of substances, 
nearly all of which, however, consist of the three ele- 
ments carbon, hydrogen, and oxygen. One of the chief 
constituents of all kinds of wood is cellulose. Now, when 
we set fire to a piece of wood, — that is to say, when we 
heat it up to the temperature at which oxygen begins to 
act on it, — it burns, if air is present. Under ordinary 
circumstances the chemical changes that take place are 
complex; but if care is taken, the combustion can be made 
complete, when all the carbon is converted into carbon 
dioxide, and all the hydrogen into water. If, on the other 
hand, the air is prevented from coming in contact with 
the wood, as by heating it in a closed vessel, or if it is 
prevented from coming in contact with it sufficiently to 
effect complete combustion, the hydrogen is given off 
partly as water and partly in the form of volatile products 



AMORPHOUS CARBON. 349 

containing carbon and oxygen, as wood spirits or methyl 
alcohol, pyroligneous or acetic acid, acetone, etc. The 
carbon, however, is for the most part left behind as char- 
coal, as there is not enough oxygen to convert it into 
carbon dioxide. Such a process as that just described, 
when carried on in closed vessels, is known as destructive 
distillation or dry distillation. It is also known as the 
charring process. It is a complex example of a kind ol 
change which we have already had to deal with. When- 
ever chemical compounds are heated the constituents tend 
to arrange themselves in forms that are stable at the higher 
temperature. Sulphites become sulphates; phosphites 
become phosphates; chlorates become perchlorates ; am- 
monium salts break down into the acids and ammonia; 
ammonium nitrite is decomposed into nitrogen and water ; 
ammonium nitrate yields nitrous oxide and water; primary 
phosphates yield metaphosphates; secondary phosphates 
yield pyrophosphates, etc., etc. Carbon compounds are, 
in general, more sensitive to the influence of heat than 
the compounds of other elements, and all are decomposed 
even at comparatively low temperatures. 

The above statements will make it possible to under- 
stand the working of a charcoal-kiln. This consists 
essentially of a pile of wood arranged to leave spaces 
between the pieces, and covered with some rough material 
through which the air will not pass easily, as, for example, 
a mixture of powdered charcoal, turf, and earth. Small 
openings are left in this covering, so that after the wood 
is kindled it will continue to burn slowly. The process 
is sometimes carried on in structures of brick-work with 
the necessary number of small openings in the walls. The 
changes above mentioned take place, the gases or volatile 
substances passing out of the top of the kiln, and appear- 
ing as a dense cloud. In due time the holes through 
which the air gains access to the wood, thus making the 
burning possible, are closed, and the burning ceases. 
Charcoal, which is impure amorphous carbon, is left 
behind. As wood always contains some incombustible 



35° COLLEGE CHEMISTRY. 

substances in small quantity, these are, of course, found 
in the charcoal. When the wood or charcoal is burned, 
these substances remain behind as the ash. 

Ordinary charcoal is a black, comparatively soft sub- 
stance. It burns in the air, though not easily, unless the 
gases that are formed are constantly removed and fresh 
air is supplied, — conditions which are brought about by a 
good draught, or by blowing upon the fire with a bellows. 
It burns readily in oxygen. The product of the combus- 
tion in the air and in oxygen, when the conditions are 
favorable, is carbon dioxide, C0 2 . In the air, when the 
draught is bad, another compound of carbon and oxygen, 
carbon monoxide, CO, is formed. Heated without access 
of air, charcoal remains unchanged. Charcoal is insoluble 
in liquids generally, though it is soluble in molten iron, 
and it crystallizes from the solution, as we have seen, in 
the form of graphite, and sometimes as diamond. 

Coke. — Besides wood charcoal, there are other forms of 
amorphous carbon, which are manufactured for special 
purposes, or are formed in processes carried on for the 
sake of other products. Coke is a form of amorphous 
carbon which is made by heating ordinary gas-coal without 
access of air, as is done on the large scale in the manufac- 
ture of illuminating gas. Coke bears to coal much the 
same relation that charcoal bears to wood. 

Lamp-black is a very finely-divided form of charcoal 
which is deposited on cold objects placed in the flames of 
burning oils. The oils consist almost exclusively of carbon 
and hydrogen. When burned in the air they yield carbon 
dioxide and water. If the flame is cooled down by any 
means, or if the supply of air is partly cut off, the carbon 
is not completely burned, the flame " smokes," as we say, 
and deposits soot. This process is chemically analogous 
to the deposit of metallic arsenic from a flame of arsine. 
The soot obtained from the flames of burning oils is made 
up largely of fine particles of carbon, though some of the 
unchanged oils are contained in it. It is used in the 
manufacture of printing-ink. As carbon is acted upon 






AMORPHOUS CARBON. 35 * 

directly by very few substances, and is not soluble, it is 
almost impossible to destroy the color of printing-ink 
without destroying the material upon which it is im- 
pressed. 

Bone-black, or Animal Charcoal, is a form of amorphous 
carbon that is made by charring bones. Bones consist of 
about one-third organic matter and two-thirds incom- 
bustible matter, mostly calcium phosphate. When charred, 
the organic matter undergoes the changes briefly described 
under the head of Charcoal, while the incombustible con- 
stituents remain unchanged. As the organic matter is 
distributed through the substance of the bones the char- 
coal obtained in this way is in a very fine state of division, 
but it is mixed with several times its own weight of min- 
eral matter. In order to remove the latter the bone-black 
must be treated with an acid, as hydrochloric acid, and 
afterwards thoroughly washed with water. An efficient 
variety of animal charcoal is made, further, by mixing 
blood with sodium carbonate, charring, and afterwards 
dissolving out the sodium carbonate with water. 

Bone-black and wood -charcoal are very porous, and 
have the power to absorb gases. When placed in air 
containing bad-smelling gases these are absorbed, and the 
air is thus to some extent purified. When water contain- 
ing disagreeable substances is treated with charcoal, these 
are wholly or partly absorbed, and the water is improved. 
Charcoal-filters are therefore extensively used. A char- 
coal-filter to be efficient should be of good size, and from 
time to time the charcoal should be taken out and renewed. 
The small filters which are screwed into faucets are of little 
value, as the charcoal soon becomes charged with the 
objectionable material which is present in the water, and 
is then a source of contamination rather than a means of 
purification. The power of charcoal to absorb gases 
depends upon its porosity. That from some varieties of 
wood is more porous than that from other varieties. Box- 
wood charcoal has been shown to absorb 90 times its own 
volume of ammonia gas, 35 times its volume of carbon 



35 2 COLLEGE CHEMISTRY. 

dioxide, and nearly twice its volume of hydrogen. Char- 
coal from cocoa-nut wood absorbs 172 times its volume of 
ammonia, and 68 times its volume of carbon dioxide. 

Some coloring matters can be removed from liquids by 
passing the liquids through bone-Mack filters. On the 
large scale, this fact is taken advantage of in the refining 
of sugar. The solution of sugar first obtained from the 
cane or beet is highly colored; and, if it were evaporated, 
the sugar deposited from it would be dark-colored. If, 
however, the solution is first passed through bone-black 
filters, the color is removed, and now, on evaporating, 
white sugar is deposited. In the laboratory constant use 
is made of this method of decolorizing liquids. The 
action can easily be shown by adding a little bone-black 
to a solution containing some litmus or indigo. If the 
solution is digested for a short time with the bone-black, 
and then passed through a filter, it will be found that the 
coloring matter is removed. 

Charcoal does not undergo decay in the air or under 
water nearly as readily as wood. That is another way of 
stating the chemical fact that the substances of which 
wood is made up are more susceptible to the action of 
other chemical substances than charcoal is. We have one 
good illustration of this, indeed, in the relative ease with 
which charcoal and wood burn in the air. Piles that are 
driven below the surface of water are sometimes charred 
to protect them from the action of those substances which 
cause decay. 

Coal. — Under this head are included a great many kinds 
of impure amorphous carbon which occur in nature. 
Although we might distinguish between an almost infinite 
number of kinds of coal, for ordinary purposes they are 
divided into hard axidi soft coals, or anthracite and bitumin- 
ous coals. Then there are substances more nearly allied 
to wood called lignite, and those which represent a very 
early stage in the process of coal-formation, viz., peat. 
A close examination of all these varieties has shown that 
they have been formed by the gradual decomposition of 



DIAMOND, GRAPHITE, AND CHARCOAL. 353 

vegetable matter in an insufficient supply of air. The 
process has been going on for ages. Sometimes the sub- 
stances have, at the same time, been subjected to great 
pressure, as can be seen from the position in which they 
occur in the earth. The products in the earlier stages of 
the coal-forming process are more closely allied to wood 
than those in the later stages. All forms of coal contain 
other substances in addition to the carbon. The soft 
coals are particularly rich in other substances. When 
heated they give off a mixture of gases and the vapors of 
volatile liquids. The gases are, for the most part, useful 
for illuminating purposes. The liquids form a black, 
tarry mass known as coal tar, from which many valuable 
compounds of carbon are obtained. The gases are passed 
through water for the purpose of removing certain impuri- 
ties. This water absorbs ammonia, and forms the am- 
moniacal liquor of the gas-works, which, as has been 
stated, is the principal source of ammonia. 

Diamond, Graphite, and Charcoal are Different Forms 
of the Element Carbon. — AVe have seen that oxygen 
presents itself in two forms — ordinary oxygen and ozone. 
Ozone is made from oxygen, and oxygen from ozone, with- 
out any increase or decrease in weight; and iu general the 
compounds obtained by the combination of other elements 
with oxygen are identical with those obtained by the com- 
bination of the same elements with ozone. So, also, there 
are several varieties of sulphur, two of which crystallize in 
different forms. There are, further, three or four differ- 
ent modifications of the element phosphorus, and these 
differ from one another in a very marked way. The ex- 
planation of the difference between oxygen and ozone is 
that the molecule of the former is made up of two atoms, 
while that of ozone is made up of three, which are in a 
state of unstable equilibrium. This explanation is reached 
through a study of the specific gravity of the two gases. 
At present no satisfactory explanation can be given of the 
difference between the varieties of phosphorus and between 
the varieties of sulphur. It will probably be shown to be 



354 COLLEGE CHEMISTRY. 

due to the way in which the atoms are grouped together 
in the molecules, and also to the way in which the mole- 
cules are grouped together to form the masses. Carbon, 
as we have seen, occurs in three distinct forms. It is 
difficult to conceive that the black, porous charcoal, and 
the dull, gray, soft graphite are chemically identical with 
the hard, transparent, brilliant diamond. Yet this is 
undoubtedly the case, as can be shown by a very simple 
experiment. Each of the substances when burned in 
oxygen yields carbon dioxide. Now, the composition of 
carbon dioxide is known, so that, if the weight of the 
carbon dioxide formed in a given experiment is known, 
the weight of the carbon in it is also known. When a 
gram of pure charcoal is burned it yields 3f grams carbon 
dioxide, and in this quantity of carbon dioxide there is 
contained exactly one gram of carbon. Further, when a 
gram of graphite is burned the same weight (3f grams) of 
carbon dioxide is obtained as in the case of charcoal; and 
the same thing is true of diamond. It follows from these 
facts that the three forms of matter known as charcoal, 
graphite, and diamond consist only of the element carbon. 
The explanation of the difference is not known, but, as 
in the cases of phosphorus and sulphur, it will probably 
be found to be in the different ways in which the atoms 
are arranged in the molecule, and the molecules in the 
masses. 

Notwithstanding the marked differences in their appear- 
ance and in many of their physical properties, the three 
forms of carbon have, as we have seen, some properties in 
common. They are insoluble in all known liquids at 
ordinary temperatures. They are tasteless, inodorous, and 
infusible. When heated without access of air they remain 
unchanged, unless the temperature is very high, when the 
diamond swells up and is converted into a mass resembling 
coke — a change which is connected with a rearrangement 
of the particles in an irregular way, so that the substances 
cease to be crystalline, or become amorphous. 



CHEMICAL CONDUCT OF CARBON. 355 

Chemical Conduct of Carbon. — At ordinary temperatures 
carbon is an inactive element. If it is left in contact with 
any one of the elements, no chemical change takes place. 
It will not combine with any of them unless the tempera- 
ture is raised. At higher temperatures, however, it com- 
bines with several of them with great ease, especially with 
oxygen. Under proper conditions it combines also with 
nitrogen, with hydrogen, with sulphur, and with many 
other elements. It combines with oxygen either directly, 
as when it burns in the air or in oxygen; or it abstracts 
oxygen from some of the oxides. The direct combination 
of oxygen and carbon has already been seen in the burn- 
ing of charcoal in oxygen, and is familiar to every one in 
the fire in a charcoal furnace. That carbon dioxide is the 
product formed can be shown by passing the gas through 
lime-water or baryta-water, when insoluble calcium or 
barium carbonate will be thrown down. The reason why 
lime-water or baryta-water is used is simply that an insolu- 
ble compound is formed, and this can be seen, and it can 
be separated from the liquid and examined. The reaction 
which takes place is represented thus: 

Ca(OH) 2 + C0 2 = OaC0 3 + H 2 0; 

Calcium Carbon Calcium 

hydroxide dioxide carbonate 

Ba(OH) 2 + C0 2 = BaC0 3 -f H 2 0. 

Barium Carbon Barium 

hydroxide dioxide carbonate 

No other common gas acts in this way on these solutions. 
Hence, when, under ordinary circumstances, a gas is 
passed into lime-water and an insoluble compound is 
formed, we may conclude that the gas is carbon dioxide, 
though this conclusion may require further proof. 

The abstraction of oxygen from compounds by means of 
carbon may be illustrated in a number of ways. Thus, 
when powdered copper oxide, CuO, is mixed with 
powdered charcoal, and the mixture heated in a tube, 
carbon dioxide is given off, and can be detected as in the 
last experiment mentioned. Copper is left behind, and, if 
the proportions are properly selected, all the carbon will 



35 6 COLLEGE CHEMISTRY. 

pass off as carbon dioxide, and only the copper be left 
behind : 

2CuO + C = 2Cu + C0 2 . 

In a similar way, arsenious oxide, As 2 3 , gives up its 
oxygen to carbon. This fact furnishes indeed a delicate 
method for the detection of the substance. If a little is 
placed in the bottom of a small tube, and above it a small 
piece of charcoal, then when heat is applied the arsenious 
oxide sublimes, and as its vapor passes the heated charcoal 
the oxygen is abstracted, and the element arsenic, being 
also somewhat volatile, is deposited just above the charcoal 
in the form of a lustrous mirror on the walls of the tube. 
The reaction is 

2As 2 3 + 30 = 4As + 3C0 2 . 

As has already been explained, the abstraction of oxygen 
from a compound is known as reduction. Hence, carbon 
is called a reducing agent. It is indeed the reducing 
agent which is most extensively used in the arts. Its chief 
use is in extracting metals from their ores, which are the 
forms in which they occur in nature. Thus, iron does 
not occur in nature as iron, but in combination with other 
elements, especially with oxygen. In order to get the 
metal the ore must be reduced, or, in other words, the 
oxygen must be extracted. This is invariably accomplished 
by heating it with some form of carbon, either coke or 
charcoal. 

The elements chlorine, oxygen, nitrogen, and hydrogen 
being gases, and the products formed when the first three 
combine with hydrogen being also gaseous or convertible 
into vapor, it is a comparatively easy matter to study the 
relations between the volumes of the combining gases and 
the volumes of the products formed. It is, however, im- 
possible to determine the ratio between the volume of 
carbon gas and that of other gases with which it combines. 

Compounds of Carbon with Hydrogen, or Hydrocarbons, 
— Carbon forms a large number of compounds with 
hydrogen. These so-called hydrocarbons, among which 



EXPERIMENTS WITH CARBON. 



357 



may be mentioned marsh-gas or methane, CJI 4 , olefiant 
gas or ethylene, C 2 H 4 , acetylene, C 2 H 2 , and benzene, 6 H 6 , 

will be treated of in a later chapter together with some 
other typical compounds of carbon. 



EXPERIMENTS. 
Carbon. — Bone-black Filters. 
Experiment 170.— Make a filter of bone-black by fitting a paper 
filter into a funnel 12 to 15 mm. (5 to 6 inches) in diameter at 
its mouth. Half fill this with bone-black. Pour a dilute solution 
of indigo through the filter. If the conditions are right the solu- 
tion will pass through colorless. Do the same thing with a dilute 
solution of litmus. If the color is not completely removed by one 
filtration, heat and filter again. The color can also be removed 
from solutions by putting some bone-black into them and boiling 
for a time. Try this with half a litre each of the litmus and 
indigo solutions used in the first part of the experiment. Use 
about 4 to 5 grams bone-black in each case. Shake the solution 
frequently while heating. 

Charcoal absorbs Gases. 
Experiment 171. — Collect over mercury in glass tubes some 
ammonia gas, and some carbon dioxide. Introduce into each 




Fig. C5. 

a piece of charcoal, which has been heated in a Bunsen-burner 
flame in order to drive out gases which may be contained in the 
pores. 



35 8 COLLEGE CHEMISTRY. 

Carbon combines with Oxygen to form Carbon 
Dioxide. 

Experiment 172. — Put a small piece of charcoal in a piece 
of hard-glass tube. Heat the tube, and pass oxygen through it. 
Pass the gases into clear lime-water. Arrange the apparatus as 
shown in Fig. 65. 

J. is a large bottle containing oxygen ; B is a cylinder contain- 
ing sulphuric acid ; C is a U-tube containing calcium chloride ; 
J) is the hard-glass tube containing the charcoal ; E is the cylin- 
der with clear lime-water. Explain all that takes place. 

Carbon reduces some Oxides when heated with them. 

Experiment 173. — In a small hard-glass tube closed at one end 

put a little of a mixture of about equal parts of powdered copper 

oxide and powdered charcoal. Attach a 

short piece of rubber tubing to the open 

end of the glass tube, as shown in Fig. 

66. Heat the tube and pass the gas which 

is given off into lime-water contained in 

a small test-tube. Is it carbon dioxide ? 

What evidence have you that oxygen has 

been extracted from the copper oxide ? 

Compare the substance left in the tube 

with metallic copper. Treat both with 

Fig. 66 nitric acid, w T ith sulphuric acid. 

Experiment 174. — Repeat Experiment 165 with somewhat 

larger quantities of the substances, and examine the gas given 

off. 

Hydrocarbons. 

Experiment 175. — Make marsh-gas by heating in a retort a 
mixture of 20 grams sodium acetate, 20 grams potassium hydrox- 
ide, and 30 grams slaked lime. Collect some of the gas over 
water. Is it a combustible gas ? 

Experiment 176. — Make ethylene as follows : In a flask of 2 to 
3 litres capacity put a mixture of 25 grams alcohol and 150 grams 
ordinary concentrated sulphuric acid. Heat to 160° to 170°, and 
add gradually through a funnel-tube about 500 cc. of a mixture 
of 1 part of alcohol and 2 parts of concentrated sulphuric acid. 
Pass the gas through three wash bottles containing, in order, con- 
centrated sulphuric acid, caustic soda, and concentrated sul- 
phuric acid. Collect some of the gas over water. Is it com- 
bustible ? 




CHAPTER XX. 

SIMPLER COMPOUNDS OF CARBON WITH OXYGEN, 
AND WITH OXYGEN AND HYDROGEN. 

General. — The final product of oxidation of carbon is 
carbon dioxide, and the final product of reduction is 
marsh-gas, bnt between these two limits there are a num- 
ber of interesting derivatives, just as there are a number 
of compounds of sulphur between hydrogen sulphide and 
sulphuric acid; a number of compounds of nitrogen 
between ammonia and nitric acid; and a number of com- 
pounds of phosphorus between phosphine and phosphoric 
acid. Some of these will be treated of in a later 
chapter. 

Carbon Dioxide, C0 2 . — The principal compound of 
carbon and oxygen is carbon dioxide, C0 2 , commonly 
known as carbonic acid gas. Under the head of The Air 
attention was called to the fact that this gas is a constant 
constituent of the air, though its relative quantity is small 
— about 3 parts in 10,000. It issues from the earth in 
many places, particularly in the neighborhood of volcanoes. 
Many mineral waters contain it in large quantity, promi- 
nent among which are the waters of Pyrmont, Selters, 
and the Geyser Spring of Saratoga. In small quantity it 
is present in all natural waters. In combination with 
bases it occurs in enormous quantities, particularly in the 
form of calcium carbonate, CaC0 3 , varieties of which are 
ordinary limestone, chalk, marble, and calc-spar. Dolo- 
mite, which forms mountain-ranges, being particularly 

359 



360 COLLEGE CHEMISTRY. 

abundant in the Swiss Alps, is a compound containing 
calcium carbonate and magnesium carbonate, MgC0 3 . 

Carbon dioxide is constantly formed in many natural 
processes. Thus, all animals that breathe in the air give 
off carbon dioxide from the lungs. That the gases from 
the lungs contain carbon dioxide can easily be shown by 
passing them through lime-water, when a precipitate of 
calcium carbonate is formed. 

That carbon dioxide is formed in the combustion of 
charcoal and wood has already been shown. In a similar 
way it can be shown that the gas is formed whenever any 
of our ordinary combustible substances are burned. From 
our fires, as from our lungs, and from the lungs of all 
animals, then, carbon dioxide is constantly given off. 
Further, the natural processes of decay of both vegetable 
and animal matter tend to convert the carbon of this 
matter into carbon dioxide, which then finds its way prin- 
cipally into the air. The process of alcoholic fermenta- 
tion, and some other similar processes, also give rise to the 
formation of carbon dioxide. In all fruit-juices there is 
contained sugar. When the fruits ripen, fall to the earth, 
and undergo spontaneous change, the sugar is converted 
into alcohol and carbon dioxide. We see, thus, that there 
are many important sources of supply of carbon dioxide, 
and it will be readily understood why the gas should be 
found everywhere in the air. 

Preparation.— The easiest way to get carbon dioxide 
not mixed with other substances is by adding an acid to 
a salt of carbonic acid or a carbonate. In the decomposi- 
tion of the carbonates by other acids we see exemplified 
the same principle as that which is involved in setting 
nitric acid free from a nitrate, or hydrochloric acid from 
sodium chloride by sulphuric acid, and more particularly 
in the liberation of sulphur dioxide from a sulphite. In 
all these cases the products are volatile, and therefore, 
when a non-volatile acid is added to the salts, decomposi- 
tion takes place. Sulphites do not yield the corresponding 
acid, but this breaks down into water and the anhydride : 



CARBON DIOXIDE. 3 6t 

( 2NaN0, + H 2 S0 4 = Na 2 S0 4 + 2HN0,; 
( 2NaCl ' + 1I.;S0 4 = Na 2 S0 4 + 2HC1. 

j Na 2 SO, + LI 2 S0 4 = N^SO, + H 2 SS0 3 ; 
\ H 2 S0 3 = S0 2 + H 2 0. 

j Na 2 C0 3 + H 2 S0 4 = Na 2 S0 4 + H 2 C0 3 ; 

1h 2 co 3 =co 2 + h 2 o. 

Any acid that is not volatile at the ordinary temperature 
will decompose a carbonate and cause an evolution of 
carbon dioxide. The action between sodium carbonate 
and hydrochloric acid is represented in this way : 

Na 2 C0 3 + 2HC1 = 2NaCl + C0 2 + H 2 0; 

that between nitric acid and sodium carbonate in this way: 

Na 2 C0 3 -j- 2HN0 3 = 2NaN0 3 + C0 2 + H 2 0. 

For the purpose of preparing carbon dioxide in the 
laboratory, calcium carbonate, in the form of marble or 
limestone, and- hydrochloric acid are commonly used. 
The reaction involved is represented thus: 

CaC0 3 + 2HC1 = CaCl 2 + C0 2 +H 2 0. 

The apparatus used is the same as that used in making 
hydrogen from zinc and sulphuric acid. As the gas is 
somewhat soluble in water, it is best for ordinary purposes 
to collect it by displacement of air, the vessel being placed 
with the mouth upward, as the gas is considerably heavier 
than air. 

Properties. — Carbon dioxide is a colorless gas at ordi- 
nary temperatures. When subjected to a low temperature 
and high pressure it is converted into a liquid. Liquid 
carbon dioxide is now manufactured on the large scale for 
use as a fire-extinguisher, and for the purpose of charging 
liquids with the gas. When some of the liquid is exposed 
to the air evaporation takes place so rapidly that a great 
deal of heat is absorbed, and some of the liquid becomes 
solid. The gas has a slightly acid taste and smell. It is 
not combustible, nor does it support combustion. It is 



362 COLLEGE CHEMISTRY 

not combustible for the same reason that water is not: 
because it already holds in combination all the oxj^gen it 
has the power to combine with. Before it can burn again 
it must first be decomposed. As regards the statement 
that it does not support combustion, it should be remarked 
that this is only relatively true. The compound does not 
easily give up oxygen, but to some substances it does give 
it up, and some such substances burn in it. For example, 
the element potassium, which, as we have seen, has the 
power to decompose water, has also the power to decom- 
pose carbon dioxide if heated in it to a sufficiently high 
temperature, and when the decomposition once begins, it 
proceeds with brilliancy, the act being accompanied by a 
marked evolution of heat and light. Carbon dioxide is 
much heavier than air, its specific gravity being 1.529. A 
litre of the gas under standard conditions of temperature 
and pressure weighs 1.977 grams. It dissolves in water, 
one volume of water dissolving about one volume of the 
gas at the ordinary temperature. As is the case with all 
gases, when the pressure is increased the- water dissolves 
more gas, and when the pressure is removed the gas again 
escapes. The so-called "soda-water" is simply water 
charged with carbon dioxide under pressure. The escape 
of the gas when the water is drawn is familiar to every 
one. The name soda-water has its origin in the fact that 
the carbon dioxide used in charging the water is frequently 
made from primary or acid sodium carbonate, NaHC0 3 , 
which is also called soda or bicarbonate of soda. 

Relations of Carbon Dioxide to Chemical Energy. — 
Carbon has the power to combine with oxygen, and in so 
doing a definite quantity of heat is evolved. A kilogram 
of carbon represents a certain quantity of chemical energy, 
which we can get from it first in the form of heat, and by 
transformation, in other forms of energy, as motion, elec- 
trical energy, etc. After the kilogram of carbon has been 
burned it no longer represents the energy it did in the 
form of carbon. A body of water elevated ten or fifteen 
feet represents a certain quantity of energy which can be 



RESPIRATION. 3 6 3 

obtained by allowing the water to fall upon the paddles of 
a water-wheel connected with the machinery of a mill. 
After the water has fallen, however, it no longer has the 
power to do work, or it has none of the energy which it 
possessed by virtue of its position. In order that it may 
again do work it must again be lifted. So, too, in order 
that the carbon in carbon dioxide may again do work the 
compound must be decomposed. 

Respiration. — It was stated above that carbon dioxide 
is given off from the lungs just as it is from a fire. It is 
a waste-product of the processes taking place in the animal 
body. Just as it cannot support combustion, so also it 
cannot support respiration. It is not poisonous any more 
than water is; but it cannot supply the oxygen which 
is needed for breathing purposes, and hence animals die 
when placed in it. They die by suffocation, very much 
as they do in droAvning. Any considerable increase in the 
quantity of carbon dioxide in the air above that which is 
normally present is objectionable, for the reason that it 
decreases the proportion of oxygen in the air which is 
breathed. If, however, pure carbon dioxide is introduced 
into the air, it has been found that as much as 5 per cent 
may be present without serious results to those who 
breathe it. In a badly-ventilated room in which a number 
of people are collected, and lights are burning, it is well 
known that in a short time the air becomes foul, and bad 
effects, such as headache, drowsiness, etc., are produced 
on the occupants of the room. These effects appear to be 
due, not to the carbon dioxide, but largely to other waste- 
products which are given off from the lungs in the process 
of breathing. The gases given off from the lungs consist 
of nitrogen, oxygen, carbon dioxide, and water-vapor. 
Besides these, however, there are many substances in 
small quantity, in a finely-divided condition, which contain 
carbon, and are in a state of decomposition. These act as 
poisons, and they are the chief cause of the bad effects 
experienced in breathing air which is contaminated by the 
exhalations from the lungs. As carbon dioxide is given 



364 COLLEGE CHEMISTRY. 

off from the lungs at the same time, the quantity of this 
gas present is roughly proportional to the quantity of the 
organic impurities. Hence, by determining the quantity 
of carbon dioxide it is possible to form an opinion as to 
whether the air of a room occupied by human beings is fit 
for use or not. 

As carbon dioxide is formed in the earth wherever an 
acid solution comes in contact with a carbonate, the gas is 
frequently given off from fissures in the earth. It is hence 
not infrequently found in old wells which have not been 
in use for some time, and deaths have been caused by 
descending into these wells for the purpose of repairing 
them. The gas is also frequently met with in mines, and 
is called choke-clamp by the miners. The miners are aware 
that after an explosion caused by fire-damp there is danger 
of death from choke-damp. The reason of the presence 
of this gas after an explosion is clear. When fire-damp, 
or marsh-gas, explodes with air the carbon is oxidized to 
choke-damp, or carbon dioxide, and the hydrogen to 
water. Air in which a candle will not burn is not fit for 
breathing purposes. 

Carbon Dioxide and Life. — The role played by carbon 
dioxide in nature is extremely important and interesting. 
The carbon contained in living things is obtained from 
carbon dioxide, and generally returns to this form when 
life ceases. We have seen that all living things contain 
carbon as an essential constituent. Whence comes this 
carbon ? Animals eat either the products of plant-life or 
other animals which derive their sustenance from the 
vegetable kingdom. The food of animals comes, then, 
either directly or indirectly from plants. But plants 
derive their sustenance largely from the carbon dioxide of 
the air. The plants have the power to decompose the gas 
with tho aid of the direct light of the sun, and they then 
build up the complex compounds of carbon which form 
their tissues, using for this purpose the carbon of the 
earbon dioxide which they decompose. Many of these 
compounds are fit for food for animals ; that is to say, they 



ENERGY STORED UP IN PLANTS. 365 

are of such composition that the forces at work in the 
animal body are capable of transforming them into animal 
tissues, or of oxidizing them, and thus keeping the tem- 
perature of the body up to the necessary point. That part 
of the food which undergoes oxidation in the body plays 
the same part as fuel in a stove. It is burned up with an 
evolution of heat, the carbon being converted into carbon 
dioxide, which is given off from the lungs. From fires 
and from living animals carbon dioxide is returned to the 
air, where it again serves as food for plants. When the 
life-process stops in the animal or the plant, decomposition 
begins; and the final result of this, under ordinary circum- 
stances, is the conversion of the carbon into the dioxide. 

Energy Stored up in Plants. — It will thus be seen that 
under the influence of life and sunlight carbon dioxide is 
constantly being converted into compounds containing 
carbon which are stored up in the plants. These com- 
pounds are capable of burning, and thus giving heat; or 
some of them may be used as food by animals, when they 
assume other forms under the influence of the life-process 
of the animals. As long as life continues, plants and 
animals are storehouses of energy. When death occurs, 
the carbon compounds begin to pass back to the form of 
carbon dioxide, and the chemical energy is transformed 
partly into heat, and is thus, as we say, dissipated. The 
power to do work, which the carbon compounds of plants 
and animals possess, comes from the heat of the sun. It 
takes a certain quantity of this heat, operating under 
proper conditions, to decompose a certain quantity of 
carbon dioxide, and elaborate the compounds contained in 
the plants. When these compounds are burned they give 
out the heat which was absorbed in their formation during 
the growth of the plants. These compounds are said to 
possess chemical energy. This has its origin in heat, and 
is capable of reconversion into heat. The transformation 
of the energy of the sun's heat into chemical energy lies 
at the foundation of all life. As the heat of the sun acting 
upon the great bodies of water and on the air gives rise to 



2,66 COLLEGE CHEMISTRY. 

the movements of water which are so essential to the exist- 
ence of the world as it is, so the action of the sun's rays 
on carbon dioxide, under the influence of the delicate and 
inexplicable mechanism ol the leaf of the plant, gives rise 
to those changes in the forms of combination of the ele- 
ment carbon which accompany and are fundamental to the 
wonderful process of life. 

Carbonic Acid and Carbonates. — When carbon dioxide 
is passed into water the solution has a slightly acid reac- 
tion. The solution will act upon bases and form salts. 
The formula of the sodium salt formed in this way has 
been shown to be Na 2 C0 3 ; that of the potassium salt, 
K 2 C0 3 ; etc. These salts are plainly derived from an acid, 
H 2 C0 3 , which is called carbonic acid. It is probable that 
this acid is contained in the solution of carbon dioxide in 
water. It is, however, so unstable that it breaks down 
into carbon dioxide and water: 

H.CO, = C0 2 + H 2 0. 

The formation of a salt by the action of carbon dioxide 
on a base takes place as shown in the following equations : 

2KOH + C0 2 = K 2 C0 3 + 11,0 ; 
Ca(OH) 2 + C0 2 = CaC0 3 + II~0. 

With the acid the action would take place as represented 
thus : 

2K0H + H 9 C0 3 = K 2 C0 3 -f 2II 2 ; 
Ca(OH) 2 + H 2 C0 3 = CaC0 3 + 2H 2 0. 

There is perfect analogy between the action of carbon 
dioxide and that of sulphur dioxide on basic solutions. 
With potassium hydroxide and calcium hydroxide, sulphur 
dioxide acts as represented in the following equations: 

2K0H + S0 2 = K 2 S0 3 + H 2 0; 
Ca(OH), + S0 2 = CaS0 3 + H 2 0. 



C/IRBON MONOXIDE. 3 6 7 

The products formed are sulphites or salts of sulphurous 
acid. 

Like sulphurous acid, carbonic acid is dibasic, and forms 
two series of salts, the primary and secondary, or the acid 
and normal salts. The primary or acid salts have the 
general formula IIMC0 3 , and the secondary or normal 
salts have the general formula M 2 C0 3 . Examples of the 
former are HKC0 3 , HNaC0 3 , CaH 2 (C0 3 ) 2 , etc.; and of 
the latter K 2 C0 3 , Na 2 C0 3 , CaC0 3 , BaC0 3 , etc. 

The secondary or normal salts which carbonic acid forms 
with the most strongly marked metallic elements, viz., 
potassium and sodium, are not decomposed by heat, but 
all other carbonates are decomposed by heat more or less 
easily, according to the strength of the base. Calcium 
carbonate when ignited loses carbon dioxide, and lime, or 
calcium oxide, remains behind: 

CaC0 3 = CaO + C0 2 . 

Carbon Monoxide, CO. — When a substance containing 
carbon burns in an insufficient supply of air, — as, for 
example, when the draught in a furnace is not strong 
enough to remove the products of combustion and supply 
fresh air, — the oxidation of the carbon is not complete, 
and the product, instead of being carbon dioxide, is carbon 
monoxide, CO. This compound can also be made by 
extracting oxygen from carbon dioxide. It is only neces- 
sary to pass the dioxide over heated carbon, when reaction 
takes place as represented thus : 

C0 2 + C = 2CO. 

This method of formation is illustrated in coal-fires, and 
can be well observed in an open grate. The air has free 
access to the coal, and at the surface complete oxidation 
takes place. But that part of the carbon dioxide which 
is formed at the lower part of the grate is drawn up 
through the heated coal, and is partly reduced to carbon 
monoxide. When the monoxide escapes from the upper 
part of the grate it again combines with oxygen, or burns, 



368 COLLEGE CHEMISTRY. 

giving rise to the characteristic blue flame always noticed 
above a mass of burning anthracite coal. 

The monoxide is also formed by passing steam over 
highly-heated carbon, when this reaction takes place : 

C -f H 2 = CO + H 2 . 

This is the reaction made use of in the manufacture of 
" water-gas." The gas thus obtained is largely a mixture 
of hydrogen and carbon monoxide. The gas is enriched 
by passing it through a furnace in which it is mixed with 
highly-heated vapors of hydrocarbons from petroleum. 
The main reaction, the decomposition of water by heated 
carbon, is effected in large furnaces filled with anthracite 
coal. The coal is first heated to a high temperature by 
setting fire to it, the products of combustion being 
allowed to escape. When it is hot enough, the air is 
shut oft' and steam passed rapidly in, when the decomposi- 
tion of the water by the carbon takes place. Soon the 
mass becomes so much cooled that the reaction stops. 
The steam is then cut off and air turned on again, and 
so on. 

The easiest way to make carbon monoxide is by heating 
oxalic acid, which is a compound of carbon, hydrogen, 
and oxygen, of the formula C 2 H 2 4 , with five to six times 
its weight of concentrated sulphuric acid. The change 
which takes place is represented thus : 



C,H,O t = C0 2 + CO + H,0. 



Both the dioxide and monoxide of carbon are formed. 
Both are gases. In order to separate them the mixture is 
passed through a solution of sodium hydroxide, which 
takes up the carbon dioxide, forming sodium carbonate, 
and allows the monoxide to pass. 

Carbon monoxide is a colorless, tasteless, inodorous 
gas, insoluble in water. It burns with a pale-blue flame, 
forming carbon dioxide. It is exceedingly poisonous when 
inhaled. Hence it is very important that it should not be 



EXPERIMENTS IV/TH CARBON DIOXIDE. 3 6 9 

allowed to escape into rooms occupied by human beings. 
We not infrequently hear of deaths caused by the gases 
from coal-stoves. The mo. a dangerous of the gases given 
off from these stoves is probably carbon monoxide. A 
pan of smouldering charcoal gives off this gas, and the fact 
that it is poisonous is well known. It has been used to a 
considerable extent for the purpose of suicide. The 
poisonous character of carbon monoxide has led to a great 
deal of discussion and to some legislation on the subject 
of '''water-gas." The question has been repeatedly raised 
whether the government should allow the manufacture of 
the gas. There is no doubt of the fact that it is a danger- 
ous substance, and that it should not be allowed to escape 
into the air is obvious. Wherever it is used special precau- 
tions should be taken to guard against leaking. There is 
no doubt that it is more poisonous than coal-gas. 

At high temperatures carbon monoxide has a very strong 
tendency to combine with oxygen, and is hence a good 
reducing agent. In the reduction of iron from its ores, 
the carbon monoxide formed in the blast-furnace plays an 
important part in the reducing process. At ordinary 
temperatures the gas does not combine readily with 
oxygen. When passed over some substances which are rich 
in oxygen, as, for example, chromic anhydride, Cr0 3 , and 
potassium permanganate, KMn0 4 , in acid solution, it 
takes up oxygen even at the ordinary temperature. It 
unites with chlorine in the direct sunlight, and forms the 
compound known as carbonyl chloride or phosgene, COCl 2 . 

EXPERIMENTS. 

Carbon Dioxide is formed when a Carbonate is 
treated with an acid. 

Experiment 177.— In test-tubes add successively dilute hy- 
drochloric, sulphuric, nitric, and acetic acids to a little sodium 
carbonate. In each case pass the gas given off through lime- 
water, and insert a burning stick in the upper part of each tube. 
Perform the same experiments with small pieces of marble. 



37° 



COLLEGE CHEMISTRY. 



Preparation and Properties of Carbon Dioxide. 



Experiment 178. 




Fig. 



—Arrange an apparatus as shown in Fig. 67. 
In the flask put some pieces of marble or lime 
stone, and pour ordinary hydrochloric acid 
on it. The gas should be collected by dis- 
placement of air, the vessel being placed 
with the mouth upward. Collect several 
cylinders or bottles full of the gas. Into one 
introduce successively a lighted candle, a 
burning stick, a bit of burning phosphorus. 
With another proceed as if pouring water 
from it. Pour the invisible gas upon the 
flame of a burning candle. Pour some of 
the gas from one vessel to another, and 
show that it has been transferred. Balance 
a beaker on a good-sized pair of scales, and 
pour carbon dioxide into it. If the balance 
is at all sensitive, the pan on which the 
beaker is placed will sink. 



Carbon Dioxide is given off from the Lungs. 

Experiment 179.— Force the gases from the lungs through 
some lime-water by means of an 
apparatus arranged as shown in 
Fig. 68. 

Formation of Carbonates. 

Experiment 180.— Pass carbon di- 
oxide into a solution of potassium 
hydroxide to saturation. Deter- 
mine whether a carbonate is in solu- 
tion or not. 

Experiment 181. — Pass carbon di- 
oxide into 50 to 100 cc. clear lime- 
water. Filter off the white insoluble 
substance. Try the action of a little 
acid on it. What evidence have 
you that it is a carbonate ? 

Experiment 182.— Pass carbon dioxide first through a little 
water to wash it, and then into 50 to 100 cc. clear dilute lime- 
water. Continue to pass the gas for some time after the pre- 
cipitate is formed. The precipitate dissolves. Heat the solution. 
What happens ? Explain these reactions. 




Fig. 68. 



EXPERIMENTS WITH CARBON MONOXIDE 31 l 

Preparation and Properties of Carbon Monoxide. 

Experiment 183. — Put 10 grams crystallized oxalic acid and 
50 to 60 grains concentrated sulphuric acid in an appropriate 
flask. Connect with two Woulffs flasks containing a solution of 
caustic soda. Heat the contents of the flask gently. Collect 
some of the gas over water. Set fire to some, and notice the 
characteristic blue flame. 

Carbon Monoxide is a good Reducing Agent. 

Experiment 184. — Pass carbon monoxide over some heated 
copper oxide contained in a hard-glass tube. Is the oxide re- 
duced ? How do you know ? Is carbon dioxide formed ? What 
evidence have you ? Was the carbon mouoxide used free of car- 
carbon dioxide ? If not, what evidence have you that carbon 
dioxide is formed in this experiment? 

Experiment 185. — Pass carbon dioxide over heated charcoal in 
a hard-glass tube. What is formed ? 



CHAPTER XXI. 

ILLUMINATION.— FLAME.— BLOWPIPE. 

COMPOUNDS OF CARBON WITH NITROGEN AND 

SULPHUR. 

Introduction. — As the substances used for illumination 
contain carbon, and the chemical processes involved con- 
sist largely in the oxidation of the carbon of these com- 
pounds, this is an appropriate place to treat briefly of the 
subject of illumination from a chemical point of view, as 
well as that of flame, and the blowpipe, which gives a use- 
ful form of flame much used in the laboratory. 

In all kinds of illumination, except the electric light, 
we are dependent upon flames for the light. Whether we 
use illuminating gas, a lamp, or a candle, the light comes 
from a flame. In the first case, the gas is burned directly; 
in the case of the lamp, the oil is first drawn up the wick, 
then converted into a gas, and this burns; while, finally, 
in the case of the candle, the solid material of the candle 
is first melted, then drawn up the wick, converted into 
gas, and the gas burns, forming the flame. In each case 
we have, then, to deal with a burning gas, and this burn- 
ing gas is called a flame. 

Illuminating Gas, Coal-gas. — Illuminating gas is some- 
times made from coal by heating it in closed retorts. As 
has already been explained, coal, particularly the softer 
kinds, contains compounds of carbon and hydrogen, 
together with some nitrogen and other elements. When 
it is subjected to destructive distillation, as in the manu- 
facture of coal-gas, the hydrogen passes off partly in com- 
bination with carbon, as hydrocarbons, and partly in the 

372 



ILLUMINATING GAS-FLAMES. 373 

free state. The nitrogen passes off as ammonia, and a 
large percentage of the carbon remains behind in the 
retort in the nncombined state as coke. The gases given 
off are purified, and form illuminating gas. One ton of 
coal yields on an average 10,000 cubic feet of gas. The 
value of a gas depends upon the amount of light given by 
the burning of a definite quantity. This is measured by 
comparing it with the light given by a candle burning at 
a certain rate. The standard candle is one made of 
spermaceti, which burns at the rate of 120 grains per 
hour; that is to say, a candle which, burning under ordi- 
nary conditions, loses 120 grains in one hour. The 
standard burner used for the gas is one through which five 
cubic feet of gas pass per hour. Now, to determine the 
illuminating' power of a gas, it is passed through the 
standard burner at the rate mentioned, and the light 
which it gives is compared with the light given by the 
standard candle. This comparison is easily made by 
means of an instrument called the photometer. The 
illuminating power of the gas is then stated in terms of 
the standard candle. The statement that the illuminating 
power of a gas is fourteen candles, signifies that, when 
burning at the rate of five cubic feet per hour, its flame 
gives fourteen times as much light as that of the standard 
candle. 

Flames. — Ordinarily a flame is a gas which is combining 
with oxygen. The hydrogen flame is simply the phenome- 
non accompanying the act of combination of the two gases 
hydrogen and oxygen. Owing to the fact that we are 
surrounded by oxygen, we speak of hydrogen as the burn- 
ing gas. How would it be if we were surrounded by an 
atmosphere of hydrogen ? Plainly, oxygen would then be 
a burning gas. If we allow a jet of oxygen to escape into 
a vessel containing hydrogen, a flame will appear where 
the oxygen escapes from the jet, if a light is applied. 
This is an experiment which requires special precautions, 
and, as the principle can be illustrated as well by means 
of illuminating gas, this may be used instead. Just as 



374 COLLEGE CHEMISTRY. 

illuminating gas burns in an atmosphere of oxygen, so 
oxygen burns in an atmosphere of illuminating gas. 

Kindling Temperature of Gases. — In studying the 
action of oxygen upon other substances, we learn eel that 
it is necessary that each of these substances should be 
raised to a certain temperature before it will combine \\i:a 
oxygen. This statement is as true of gases as of otner 
substances. When a current of hydrogen is allowed to 
escape into the air, or into oxygen, no action takes place 
unless it is heated up to its burning temperature, when it 
takes fire and continues to burn, as the burning of one 
part of the gas heats the part which follows it, and hence 
it is heated to the burning temperature as fast as it escapes 
into the air. If the gas is cooled even very slightly below 
this temperature, it is extinguished. This can easily be 
shown by bringing down upon the flame of a Bunsen 
burner a piece of wire gauze. There will be no flame 
above the gauze, but gas will pass through unburned, and 
this will burn if it is lighted above the gauze. In this 
case, by simply passing through the thin wire gauze, the 
gases are cooled below their burning temperatures, and 
the flame does not pass through. So, also, if the gas is 
turned on and not lighted, and the gauze held an inch or 
two above the outlet, the gas will burn above the gauze if 
lighted above, and will not pass downward through the 
gauze, unless this becomes very hot. 

Miner's Safety-lamp. — The principle illustrated in the 
experiments referred to in the last paragraph is utilized in 
the miner's safety -I amp, to which reference has already 
been made. One of the dangers which the coal-miner has 
to encounter is the occurrence in the mines of fire-damp, 
or methane, CH 4 , which with air forms an explosive mix- 
ture. The explosion can only be brought about by contact 
of flame with the mixture. In order to avoid the contact, 
the flame of the safety-lamp is surrounded by wire gauze, 
as shown in Fig. 69. When a lamp of this kind is 
brought into an explosive mixture of marsh-gas and air, 
the mixture passes through the wire gauze and comes in 



STRUCTURE OF FLAMES. 



275 




Fig. 69. 



eontaci with the flame, and a small explosion or a series 
of small explosions inside the gauze occurs, but the flame 
of the burning gas inside the wire gauze 
cannot pass through and raise the tem- 
perature of the gas outside to the burning 
temperature. Hence no serious explosion 
can take place. The flickering of the 
flame of the lamp, and the occurrence of 
small explosions inside, furnish the miner 
with the information that he is in a dan- 
gerous atmosphere. While the safety-lamp 
does undoubtedly afford much protection, 
still explosions occur. These have been 
shown to be caused by the presence of 
coal-dust in the mines, and by the com- 
motion of the air produced in blasting. 
By the aid of the coal-dust, and by sudden 
and violent movements of the air, it is possible for a flame 
surrounded by wire gauze to explode a mixture of marsh- 
gas and air on the other side of the gauze. 

Structure of Flames. — The hydrogen flame consists of 
a thin envelope of burning hydrogen enclosing unburned 
gas, and surrounded by water-vapor, which is the product 
of the combustion. The structure of other flames depends 
upon the complexity of the gases burned, and the condi- 
tions under w T hich the burning takes place, In general, 
a flame consists of an outer envelope of gas combining 
with oxygen, and hence hot, and an inner part which 
contains unburned gas, which is comparatively cool. A 
part of the unburned gas is, however, quite hot, and it 
would combine with oxygen were it not for the fact that 
it is surrounded by an envelope that prevents access of 
air. The outer hot part of the flame is called the oxidiz- 
ing flame, because it presents conditions favorable to the 
oxidation of substances introduced into it. The inner 
hot part is called the reducing flame, because it consists of 
highly-heated substances which have the power to combine 
with oxygen; and hence many compounds containing 



37^ 



COLLEGE CHEMISTRY. 



oxygen lose it, or are reduced, when introduced into this 
part of the flame. The hottest part of the name is about 
half-way between the bottom and the top. Here oxidation 
is taking place most energetically. The hottest part of 
the unburned gases is at the tip of the dark central part 
of the flame. In the flame of a Bunsen burner the two 
parts can be easily distinguished. The dark central part 
of the flame extends for some distance above the outlet of 
the burner. If the holes at the base of the burner are 
partly closed, the tip of the central part of the flame 
becomes luminous. This luminous tip is most efficient 
' for the purpose of reduction. The prin- 
2 cipal parts of the flame are those marked in 
Fig. 70. The part marked b is the central 
cone of unburned gases; that marked c is 
the luminous tip, the best part of the flame 
for reduction. A is the envelope of burn- 
ing gas. The hottest part of the flame is 
at a; that which is most efficient in causing 
oxidation is at d. This is further sur- 
rounded by a non-luminous envelope con- 
sisting of the products of combustion, 
carbon dioxide and water-vapor. Certain 
metals placed in the upper end of the flame 
take up oxygen, because they are highly heated in the 
presence of oxygen. Certain oxides lose their oxygen 
when placed in the tip of the central cone, because the 
gases are here heated to the temperature at which they 
have the power to combine with oxygen. 

Blowpipe. — The oxidizing and reducing flames are fre- 
quently utilized in the laboratory. For the purpose of 
increasing their efficiency a blowpipe is used. This is 
a tube with a convenient mouth-piece and a nozzle with 
a small opening through which air is blown into a flame. 
The blowpipe may be used with the flame of a candle, 
an alcohol-lamp, or a gas-lamp. It is commonly used 
with a gas-lamp. By regulating the current of air and 
slightly changing the position of the tip of the blow- 




Fig. 70. 



CAUSES OF THE LUMINOSITY OF FLAMES. 377 

pipe a good oxidizing flame or a good reducing flame can 

be produced. Some oxides are very easily reduced when 
heated in the reducing blowpipe flame. Others are not. 
We can frequently judge of the composition of a substance 
by heating in the blowpipe flame, and noticing its conduct. 
Some metals are easily oxidized in the oxidizing flame. 
Some form characteristic films, or thin layers of oxides, 
on the substance upon which they are heated, which is 
usually charcoal; and, in some cases, it is possible to 
detect the presence of certain substances by the color of the 
film of oxide. The blowpipe is therefore of much value 
as affording a method for the detection of the presence of 
certain elements in mixtures or compounds of unknown 
composition. The chemical principles involved in its use 
will be clear from what has already been said. 

Causes of the Luminosity of Flames. — It is evident 
from what we have seen that flames differ greatly in their 
light-giving power. The hydrogen flame, for example, 
though extremely hot, gives practically no light. This is 
also the case with the flame of the Bunsen burner; while, 
on the other hand, the flame of coal-gas, burning under 
ordinary circumstances, and that of a candle, etc., give 
light. To what is the difference due ? This subject has 
been studied very thoroughly, and it has been found that 
there are several causes that operate to make a flame 
luminous and vice versa. In the first place, if a solid sub- 
stance which does not burn is introduced into a non- 
luminous flame, a part of the heat appears as light. This 
is seen when a spiral of platinum wire is introduced into 
a hydrogen flame. It is also seen when a piece of lime is 
introduced into the hot non-luminous flame of the oxy- 
hydrogen blowpipe. A similar cause operates in ordinary 
gas-flames to make them luminous. Particles of unburned 
carbon are always present, as can be shown by putting a 
piece of porcelain or any solid substance into the flame, 
when there will be deposited on it a layer of soot, which 
consists mainly of finely-divided carbon. In the flame 
such particles are heated to incandescence, or to the tern- 



37* COLLEGE CHEMISTRY. 

perature at which they give light. Again, it has been 
found that a candle gives more light at the level of the sea 
than it does when at the top of a high mountain, as Mount 
Blanc, on which the experiment was actually performed. 
This is partly due to a difference in the density of the 
gases. Naturally, the denser the gas the more active the 
combustion, the greater the heat, and the brighter the 
light. This last statement ceases to be true when the 
oxidation becomes sufficient to burn up all the solid 
particles in the flame. If gases, which in burning give 
light, are cooled down before they are burned, the 
luminosity is diminished, and, conversely, non-luminous 
flames may be rendered luminous by heating the gases 
before burning them. Gases which otherwise give lumin- 
ous flames give non-luminous flames when diluted to a 
sufficient extent with neutral gases, such as nitrogen and 
carbon dioxide, which neither burn nor support combus- 
tion. Lewes has recently shown that acetylene (which see) 
is always formed in flames and his work makes it appear 
probable that the luminosity of most, if not of all, flames 
is due to the formation and decomposition of this gas. 

Bunsen Burner. — All the statements made in regard to 
the causes of the luminosity of flames are based upon 
carefully-performed experiments. These experiments, 
however, cannot, for the most part, be readily repeated by 
the student in the laboratory in a satisfactory way. One 
constant reminder of the possibility of rendering a 
luminous flame non-luminous, and vice versa, is furnished 
by the burner universally used in chemical laboratories, 
and called, after the inventor, the Bunsen burner. The 
construction of this burner is easily understood. It con- 
sists of a base and an upper tube. The base is connected 
by means of a rubber tube with the gas-supply. The gas 
escapes from a small opening in the base, and passes 
upward through the tube. At the lower part of the tube 
there are two holes, which may be opened or closed by 
turning a ring with two corresponding holes in it. When 
the gas is turned on, it is lighted at the top of the tube. 



CYANOGEN. 379 

Air is at the same time drawn through the holes at the 
base. The result is that the flame is practically non- 
luminous. If the ring at the base is turned so that the 
air-holes are closed, the flame becomes luminous. The 
advantage of the non-luminous flame for laboratory use 
consists in the fact that it does not deposit soot, and, at 
the same time, it is hot. 

The non-luminosity of the flame of the Bunsen burner 
appears to be due to several causes: (1) Dilution of the 
gases by means of the nitrogen of the air; (2) Cooling of 
the gases by the entrance of the air; (3) Burning of the 
solid particles by the aid of the oxygen of the air admitted 
to the interior of the flame. 



Compounds of Carbon with Nitrogen and with 
Sulphur. 

Cyanogen, C 2 N 2 . — Carbon does not combine with nitro- 
gen under ordinary circumstances. If, however, these 
elements are brought together at very high temperatures 
in the presence of metals, they combine to form com- 
pounds known as cyanides. Thus, when nitrogen is passed 
over a highly-heated mixture of carbon and potassium 
carbonate, potassium cyanide, KCN, is formed. Carbon 
containing nitrogen, as animal charcoal, when ignited 
with potassium carbonate, reduces the carbonate, forming 
potassium, in presence of which carbon and nitrogen com- 
bine, forming potassium cyanide. When refuse animal 
substances, such as blood, horns, claws, hair, etc., are 
heated together with potassium carbonate and iron, a sub- 
stance known as potassium ferro cyanide, or yellow prussiate 
of potash, K 4 Fe(CN) 6 + 3H 2 0, is formed. When this is 
simply heated it is decomposed, yielding potassium cyanide: 

KFe(CN) 6 = 4KCN + FeC 2 + N,. 

It is an easy matter to make the mercury salt, Hg(CN) 2 , 



380 COLLEGE CHEMISTRY. 

from the potassium salt. By heating mercuric cyanide it 
breaks up, yielding metallic mercury and cyanogen gas : 

Hg(CN), = Hg + C,N„ 

just as mercuric oxide yields mercury and oxygen when 
heated : 

HgO = Hg + 0. 

Cyanogen (from Kvaros, Hue) owes its name to the fact 
that several of its compounds have a blue color. It is a 
colorless gas, which is easily soluble in water and alcohol, 
and is extremely poisonous. It burns with a purple-colored 
flame. 

Hydrocyanic Acid, Prussic Acid, HON. — This acid, 
which is commonly called prussic acid, occurs in nature 
in amygdalin, in combination with other substances, in 
bitter almonds, the leaves of the cherry, laurel, etc. It is 
prepared by decomposing metallic cyanides with hydro- 
chloric acid. It is volatile and passes over. The action 
is represented thus : 

KCN + HOI = KC1 + HCN. 

Hydrocyanic acid is a volatile liquid, boiling at 26.1°, 
and solidifying at — 14°. It has a characteristic odor 
suggestive of bitter almonds. It is extremely %)oisonous. 
It dissolves in water in all proportions. Such a solution 
is known as prussic acid. Pure hydrocyanic acid is very 
unstable. By standing, a brown substance is deposited 
from its solution. 

Cyanic Acid, IICNO. — By gentle oxidation of a cyanide 
it is converted into a cyanate. Thus, by melting together 
potassium cyanide and lead oxide, potassium cyanate is 
formed : 

KCN + PbO = KCNO + Pb. 

Cyanic acid is a volatile, acrid, unstable liquid. It 



CARBON BISULPHIDE. 381 

breaks down at once into carbon dioxide and ammonia in 
presence of water : 

CONH + H 2 = NH 3 + C0 2 . 

The potassium salt is easily soluble in water, but is 
decomposed by it, yielding ammonia and acid potassium 
carbonate : 

CONK + 2H 2 = KHCO3 + NH 3 . 

These decompositions of cyanic acid and the cyanates 
further exemplify the tendency of cyanogen compounds 
to undergo decomposition in presence of water. 

Carbon Disulphide, CS 2 . — Just as carbon combines 
directly with oxygen to form the dioxide, so it combines 
directly with sulphur to form the disulphide; but there is 
a great difference in the ease with which carbon combines 
with the two elements. In order to effect combination 
with sulphur a very high temperature is necessary. The 
compound is prepared on the large scale by heating char- 
coal to a high temperature in an upright cast-iron cylinder, 
and adding sulphur in such a way that it enters the bottom 
of the cylinder. The product is passed through a series 
of tubes arranged so as to secure condensation. 

Carbon disulphide is a clear liquid which has a high 
refractive power. It boils at 40.2°. When pure it has a 
pleasant odor, but if kept for a time, particularly if water 
is present in the vessel, it undergoes slight decomposition, 
and products of extremely disagreeable odor are formed. 
It can generally be freed from these by shaking it with a 
little mercury and then redistilling. It burns readily, 
forming carbon dioxide and sulphur dioxide : 

CS 2 + 30 2 = CO, + 2S0 2 . 

Carbon disulphide is only slightly soluble in water, and 
is decomposed by it only very slowly. The disulphide is 



3^2 COLLEGE CHEMISTRY. 

an excellent solvent for many substances which are not 
soluble in water, as> for example, fats, resins, iodine, and 
one of the modifications of sulphur and of phosphorus. 
The solution of iodine in it has a beautiful violet color; 
and when a water solution containing a little free iodine 
is shaken with carbon disulphide the latter acquires a 
violet color and separates below the water. 

Carbon disulphide finds extensive application as a 
solvent, and it is also used for the purpose of destroying 
phylloxera, the insect which is so destructive to grape-vines, 
particularly in the wine districts of France. 

Sulphocyanic Acid, HCNS. — Just as the cyanides take 
up oxygen and are converted into cyanates, so also they 
take up sulphur and are converted into sulphocyanates : 

KCN + S = KCNS. 

Constitution of Cyanogen and its Simpler Compounds. 

— The compounds of cyanogen show, in general, a remark- 
able similarity to the compounds of the chlorine group. 
The hydrogen compound is a monobasic acid and forms a 
series of salts, the cyanides, which in general are analogous 
to the chlorides. Comparing the cyanides with the 
chlorides it is clear that in the former the group (CN), or 
the cyanogen group, plays the same part that the atom 
chlorine plays in the chlorides : 

H(CN) HC1 

K(CN) KC1 
Hg(CN) 2 HgCl, 

So, also, cyanic acid and hypochlorous acid are analogous: 

HO(CN) IIOC1. 

This relation suggests that which is observed between the 
ammonium compounds and those of potassium and sodium. 
The cyanogen group is evidently univalent, as it combines 
with one atom of hydrogen, one of potassium, etc. 



EXPERIMENTS : COAL-GAS, ETC. 



3*3 



EXPERIMENTS. 

Coal-gas. 

Experiment 186. — Heat some bituminous coal in a retort and 
collect over water the gases given off. Are these gases com- 
bustible ? 

Oxygen burns in an Atmosphere of a Combustible 

Gas. 

Experiment 187. — Break off the neck of a good-sized retort ; 
fit a perforated cork to the small end ; pass a piece of glass tube 
through the cork, and connect by means of rubber hose with an 
outlet for coal-gas. Fix the apparatus in position, as shown in 
Fig. 71. Turn the gas on, and when the air is driven out of the 




Fig. 71. 

retort-neck, light the gas. The neck is now filled with illu- 
minating gas, and the gas is burning at the mouth of the vessel. 
If now a platinum jet from which oxygen is issuing is passed up 
into the gas the oxygen will take fire, and a flame will appear 
where the oxygen escapes from the jet. The oxygen " burns " in 
the atmosphere of coal-gas. 

Kindling Temperature of Gases. 

Experiment 188. — Light a Bunsen burner. Bring down upon 
the flame a piece of brass or iron wire-gauze. There is no flame 
above the gauze. That the gas passes through unburned can be 
shown by applying a light just above the outlet of the burner 



3§4 



COLLEGE CHEMISTRY. 



and above the gauze. The gas will take fire and burn. By 
simply passing through the thin wire-gauze, then, the gas is 
cooled down below its burning temperature, and does not burn 
unless it is heated up again. Turn on a Bun sen 
burner. Do not light the gas. Hold a piece of wire- 
gauze about one and a half to two inches above the 
outlet. Apply a lighted match above the gauze, 
when the gas will burn above the gauze, but not 
below it. Here again the heat necessary to raise the 
temperature of the gas to the burning temperature 
cannot be communicated through the gauze. If in 
either of the above-described experiments the gauze 
is held in position for a time, it will probably become 
so highly heated that the gas on the side where there 
is no flame will be raised to the burning tempera- 
ture. The instant that point is reached the flame 
becomes continuous. 



4- 



The Blowpipe and its Uses. 



Fig. 72. The blowpipe used in chemical laboratories is con- 
structed as shown in Fig. 72. 
When used w r ith the Bunsen burner it is best to slip into the 
burner a brass tube ending above in a narrow slit-like opening, 
as shown in Fig. 73. The tube re- 
ferred too, marked a in the figure, 
reaches to the bottom of the burner, 
and thus cuts off the supply of air 
which usually enters the holes at the 
base. The gas is now lighted, and the 
current so regulated that there is a 
small flame about 1\ to 2 inches long. 
The tip of the blowpipe is placed on the 





Fig. 73. Fig. 74. 

slit of the burner in the flame, as shown in Fig. 74. By blowing 



THE BLOWPIPE AND ITS USES. 3 8 5 

regularly and not violently through the pipe the flame is forced 
down in the same direction as the end-piece of the blowpipe, and 
the slant of the burner-slit. Under proper conditions the flame 
separates sharply into a central blue part and an outer part of 
another color. The direction and lines of division of the flame 
are indicated in Fig. 74. The outer part of the flame marked o is 
the oxidizing flame ; the part marked r is the reducing flame. 

Experiment 189. — Select a piece of charcoal about 4 inches 
long by 1 inch wide and 1 inch thick, with one surface plane.* 
Near the end of the plane surface make a cavity by pressing the 
edge of a small thin coin against it, and turning it completely 
round a few times. Mix together equal small quantities of dry 
sodium carbonate and lead oxide. Put a little of the mixture in 
the cavity in the charcoal, and heat it in the reducing flame 
produced by the blowpipe. In a short time globules of metallic 
lead will be seen in the molten mass. After cooling, scrape the 
solidified substance out of the cavity in the charcoal. Put it in a 
small mortar, treat it with a little water, and, after breakiug it 
up and, allowing as much as possible to dissolve, pick out the 
metallic beads. Is it malleable or brittle ? Is metallic lead mal- 
leable or brittle ? Is it dissolved by hydrochloric acid ? Is lead 
soluble in hydrochloric acid? Is it soluble in nitric acid? Is 
lead soluble in nitric acid ? The action of the acids can be tried 
by putting a bead on a small dry watch-glass and adding a few 
drops of the acid. Does the substance act like lead ? What has 
become of the oxygen with which the lead was combined in the 
oxide ? Is there any special advantage in having a support of 
charcoal for this experiment ? 

Experiment 190. — Heat a small piece of metallic lead on char- 
coal in the oxidizing blowpipe flame. Notice the formation of 
the oxide, which forms a coating or film on the charcoal in the 
neighborhood of the metal. Is there any analogy between this 
process and the burning of hydrogen ? In what does the analogy 
consist ? What differences are there between the two processes ? 

Experiment 191. — Repeat the experiments with arsenic, anti- 
mony, and bismuth. Notice the colors of the films formed on 
the charcoal. 

Experiment 192. — Melt into a bit of glass tubing a piece of 
platinum wire 8 to 10 mm. (3 to 4 inches) long, and bend the end 
so as to form a small loop, as shown in Fig. 75. Heat the loop in 
the flame of a Bunsen burner, and then dip it into some sodium- 

* Pieces of charcoal prepared for blowpipe work can be bought 
from dealers in chemical apparatus at small cost. 



3 86 COLLEGE CHEMISTRY. 

ammonium phosphate (microcosmic salt). Heat in the oxidizing 
flame of the blowpipe until a clear glass bead is formed in the 
loop. What changes have taken place? and what is the clear 

T-r-r-" ■-■■■--■--:•■■■$ D 

Fig. 75. 

glass ? Bring a minute particle of a manganese compound in 
contact with the bead, and heat again. What change takes 
place ? Try the same experiment, using successively a cobalt 
compound, a copper compound, and an iron compound. Now, 
instead of using microcosmic salt, use borax. Explain the 
changes in all the above-described experiments. 

Cyanogen. 

Experiment 193. — Make potassium cyanide by heating potas- 
sium ferrocyanide in an iron crucible. 

Experiment 194. — Make cyanogen by heating mercuric cy- 
anide. Cyanogen is poisonous. Burn some of the gas. 

Experiment 195. — Make potassium cyanate from some of the 
cyanide obtained in Experiment 193. This is done by melting it 
in an iron crucible, and, while the mass is liquid, adding about 
four times its weight of red lead, stirring during the operation. 
After this the crucible should again be put in the furnace for a 
little while, the metallic lead allowed to settle, and the contents 
poured out on a smooth stone. Break this up, and extract the 
cyanate with alcohol. 



CHAPTER XXII. 

ELEMENTS OF FAMILY IV, GROUP A: 

SILICON.— TITANIUM.— ZIRCONIUM.— CERIUM.— 

THORIUM. 

General. — While in some respects the elements of this 
group resemble carbon and bear to it relations similar to 
those which the members of the chlorine group bear to 
fluorine, the members of the sulphur group to oxygen, and 
the members of the phosphorus group to nitrogen, yet 
between them and carbon there are some remarkable 
differences. All the members of the group except titanium 
combine with hydrogen. The compounds formed have 
the formulas SiH 4 , ZrH 2 , CeH 2 , and ThH 2 . 

All the elements of the group form oxygen compounds 
analogous to carbon dioxide. They are : 

Si0 2 , Ti0 2 , Zr0 2 , Ce0 2 , Th0 2 . 

The first three are acidic, and form salts which in com- 
position are analogous to the carbonates. These are the 
silicates, titanates, and zirconates of the general formulas 

M 2 Si0 3 , M 2 Ti0 3 , M 2 Zr0 3 . 

Cerium and thorium oxides are basic. These facts sug- 
gest the relations between the members of the phosphorus 
group. The oxides of the last two members, antimony 
and bismuth, are basic, although the oxide of antimony is 
also acidic in its conduct towards the stronger bases. 

3*7 



3^8 COLLEGE CHEMISTRY. 

The compounds of silicon are very abundant in nature; 
those of the other members of the group are rare. 

Silicon, Si (At. Wt. 28.4). 

Occurrence. — We have already seen what an exceedingly 
important part carbon plays in animate nature. It is in- 
teresting to note that silicon, which in some respects from 
a chemical point of view resembles carbon, is one of the 
most important constituents of the mineral or inorganic 
parts of the earth. It occurs chiefly in the form of the 
dioxide, Si0 2 , commonly called silica, or silicon dioxide; 
and in combination with oxygen and several of the common 
metallic elements, particularly with sodium, potassium, 
aluminium, and calcium, in the form of the silicates. 
Next to oxygen, silicon is the most abundant element in 
the earth. There are extensive mountain-ranges consist- 
ing almost entirely of the dioxide, Si0 2 , in the form 
known as quartz or quartzite. Other ranges are made up 
of silicates, which are compounds formed by the combina- 
. tion of silicon dioxide and bases. The clay of the valleys 
and river-beds also contains silicon in large quantity, while 
the sand found so abundantly on the deserts and at the 
seashore is largely silicon dioxide. 

Preparation. — Silicon does not occur in nature in the 
free state. The oxide, Si0 2 , which is most abundant in 
the form of sand, is decomposed by heating it with potas- 
sium or magnesium, and silicon is thus set free. When 
magnesium is used the action is violent, and besides the 
silicon a compound of silicon and magnesium is formed. 
Silicon has also been made by heating the oxide and carbon 
in the electric furnace, and by decomposing the chloride 
with potassium: 

SiCl 4 + 4K = Si + 4K01. 

The best way to make it is by heating together potassium 
fluosilicate, K 2 SiF 6 , sodium, and zinc: 

K 2 SiF 6 + 4Na = 4NaF + 2KF + Si. 



SILICON. 3 8 9 

At the same time the zinc melts and the silicon which 
separates dissolves in the molten zinc. On cooling, it is 
deposited from the solution in beautiful needle-shaped 
crystals, around which the zinc solidifies at a lower tem- 
perature. By treating the mass with hydrochloric acid 
the zinc is dissolved and the crystals of silicon are left 
behind. When obtained by reduction of the oxide or the 
chloride by means of potassium, it is a brown amorphous 
powder. If made by decomposition of potassium fluosili- 
cate by aluminium, it is deposited from the molten 
aluminium in crystals somewhat resembling graphite. 
Just as there are three forms of carbon, the amorphous, 
graphite, and diamond, so there are three corresponding 
forms of silicon, the amorphous brown powder, the graphi- 
toidal, and the needles. The amorphous variety is con- 
verted into crystallized silicon by continued heating at a 
high temperature. 

Amorphous silicon acts upon hydrofluoric acid, forming 
silicon tetrafluoride, SiF 4 , and setting hydrogen free: 

Si + 4HF = SiF 4 + 2H 2 . 

In this reaction it exhibits one of the properties of a base- 
forming element. Towards other acids, however, it is 
indifferent. It is not acted upon by sulphuric acid, nor 
by nitric acid, nor aqua regia. It dissolves, however, in 
potassium hydroxide, forming potassium silicate, in this 
case acting like an acid-forming element: 



Si + 2KOH + H 2 = K 2 Si0 3 + 2H 



This form of silicon also burns in the air, forming the 
dioxide. 

Crystallized silicon, on the other hand, does not burn 
in oxygen at the highest temperatures. It, however, 
reduces carbon dioxide and decomposes carbonates at a 
high temperature. It is also oxidized by a melting mix- 
ture of potassium nitrate and the hydroxide or carbonate. 
It combines with nitrogen at a high temperature. 



390 COLLEGE CHEMISTRY. 

Both the graphitoidal and needle-formed crystals of 
silicon consist of regular octahedrons. Both forms have 
a blackish-gray color and a metallic lustre. 

Silicon Hydride, SiH 4 . — This gas is obtained mixed with 
hydrogen when a compound of magnesium and silicon is 
treated with hydrochloric acid : 

Mg 2 Si + 4HC1 = SiH 4 + 2MgCl 2 . 

Thus made, it takes fire when it comes in contact with 
the air, and the act is accompanied by explosion. The 
products of its combustion are silicon dioxide and water. 
When pure it forms a colorless gas which does not take 
fire spontaneously in the air at the ordinary temperature. 
If it is diluted with hydrogen, or if it is heated, it does 
take fire. When burned in a cylinder or narrow tube, so 
that free access of air is not possible, amorphous silicon is 
deposited upon the walls of the vessel. 

Titanium, Ti (At. Wt. 48.1). — Titanium occurs in 
nature as titanium dioxide, Ti0 2 , in three distinct forms, 
known as rutile, brookite, and anatase; in combination 
with iron, as titaniferous iron which contains ferrous 
titanate, FeTi0 3 ; and in a number of iron ores and rare 
minerals. 

Zirconium, Zr (At. Wt. 90.7). — The principal form in 
which zirconium occurs in nature is as zircon, which is a 
silicate of the formula ZrSi0 4 , derived from normal silicic 
acid, Si(OH) 4 , by the substitution of a quadrivalent atom 
of zirconium for the four hydrogen atoms. 

Thorium, Th (At. Wt. 232.5).— This element occurs 
principally in the mineral thorite, which is essentially a 
silicate of thorium, ThSi0 4 , analogous to zircon. 

Cerium so much resembles the two elements lanthanum 
and didymium that, although it falls in the same group 
as silicon and resembles the elements of this group in 
some respects, it seems advisable to postpone its study until 
lanthanum and didymium are taken up. 



COMPOUNDS OF THE SILICON GROUP. 39 * 

Compounds of the Elements of the Silicon Group 
with those of the chlorine group. 

Silicon Tetrachloride, SiCl 4 . — This compound is formed 
when silicon is heated in a current of chlorine, and by 
passing a current of dry chlorine OTer a heated mixture of 
silicon dioxide and carbon. Under these latter circum- 
stances the following reaction takes place: 

SiO, + 2C + 2C1 2 = SiCl 4 + 2CO. 

The tetrachloride is a colorless liquid. It is decomposed 
by water, forming silicic acid and hydrochloric acid. The 
reaction probably takes place as represented in the follow- 
ing equation: 

SiCl 4 + 4H 2 = Si(OH) 4 + 4HC1. 

The normal acid thus formed breaks down readily, how- 
ever, forming the ordinary acid of the formula SiO(OH) 2 
or H 2 Si0 3 , corresponding to carbonic acid, H 2 C0 3 . 

Silicon Hexachloride, Si 2 Cl 6 , is formed when silicon 
tetrachloride is heated with silicon : 

3SiCl 4 + Si = 2Si 2 Cl 6 . 

When heated to a sufficiently high temperature it is 
decomposed, yielding silicon and the tetrachloride : 

2Si 2 Cl 6 = 3SiCl 4 + Si. 

Water decomposes it, forming the corresponding hydroxyl 
derivative, which loses water and forms the acid 
Si,0 2 (OH) 2 : 

Si 2 Cl 6 + 6H 2 = Si. 2 (OH) 6 + 6HC1; 
Si 2 (OH) 6 = Si 2 2 (OH) 2 + 2H 2 0. 

The product is a disilicic acid, in some respects analogous 
to disulphuric acid. 



39* COLLEGE CHEMISTRY. 

Silicon Tetrafluoride, SiF 4 . — This is one of the most 
interesting of the compounds that silicon forms with the 
members of Family VII. It is made by treating silicon 
dioxide with hydrofluoric acid. This action is secured by 
treating a mixture of silicon dioxide (sand) and calcium 
fluoride (fluor-spar) with concentrated sulphuric acid, 
when two reactions take place: 

CaF 2 + H 2 S0 4 = CaS0 4 + 2HF; 
Si0 2 4- 4HF = 2H 2 + SiF 4 . 

The tetrafluoride escapes as a colorless gas, that forms 
thick clouds in moist air on account of the action of water 
upon it. 

Water decomposes the tetrafluoride, as it does the tetra- 
chloride. The first action probably consists in the forma- 
tion of normal silicic acid and hydrofluoric acid, the 
normal acid then breaking down by loss of water and 
yielding the ordinary form of silicic acid : 

SiF 4 + 4H 2 = Si(OH) 4 + 4HF; 
Si(OH) 4 = SiO(OH) 2 + H 2 0. 

The silicic acid thus formed separates as a gelatinous mass. 
At the same time the hydrofluoric acid acts upon some of 
the silicon tetrafluoride, forming the compound fiuosilicic 
acid, which has the formula H 2 SiF 6 : 

SiF 4 + 2HF = H 2 SiF 6 . 

The complete action may be represented in one equation, 
as follows : 

3SiF 4 + 3H 2 = H 2 Si0 3 + 2H 2 SiF 6 . 

The fiuosilicic acid remains in solution in the water, and 
by treating this solution with carbonates or hydroxides of 
the metallic elements the salts known as the fluosilicates 
are obtained. The solution of the acid can be concentrated 
to a certain extent in a platinum vessel, but it breaks down 



COMPOUNDS OF THE SILICON GROUP. 393 

iuto silicon tetrafluoride and hydrofluoric acid when it 
becomes concentrated. If more potassium hydroxide than 
is required to neutralize the acid is added to the solution, 
decomposition ensues, with formation of silicic acid: 

H 2 SiF 6 + GKOH = 6KF + H 2 Si0 3 + 3H 2 0. 

By water alone, however, the acid is not decomposed, and 
the salts are fairly stable. When heated, the salts give off 
silicon tetrafluoride, and fluorides are left behind: 

K a SiF 6 = 2KF + SiF 4 . 



Compounds of the Members of the Silicon Group 
with Oxygen, and with Oxygen and Hydrogen. 

Silicon Dioxide, Si0 2 . — This compound occurs very 
abundantly in nature in many different forms, both 
crystallized and amorphous. Quartz is a form of crystal- 
lized silicon dioxide which is found very widely distributed. 
It crystallizes in the hexagonal system in prisms and 
pyramids, the crystals sometimes attaining great size and 
beauty. Another form of the crsytallized compound is 
that known as tridymite. Like quartz it crystallizes in 
the hexagonal system, but the characteristic forms are not 
the same as those of quartz. Further, it nearly always 
occurs in triplet crystals. The finer crystals of quartz are 
generally called rock-crystal ; the crystalline variety in 
which the crystals are not well developed is called quartz- 
ite. The amorphous varieties of silicon dioxide frequently 
contain water in combination, or, rather, they are hy- 
droxides of silicon. Examples of these forms are opal, 
agate, amethyst, carnelian, flint, sand, chalcedony Some 
of these are colored by small quantities of other substances 
contained in them. Carnelian ower its color to a com- 
pound of iron, probably ferric oxide ; flint contains small 
quantities of organic matter. The specific gravity of the 



394 COLLEGE CHEMISTRY. 

crystallized varieties is higher than that of the amorphous 
varieties, and there are also some chemical differences 
between them. 

Properties. — Silicon dioxide is insoluble in water and in 
most acids. It dissolves, however, in hydrofluoric acid, 
forming the tetrafluoride. It requires the temperature 
produced by the oxyhydrogen blowpipe to melt it. The 
amorphous varieties are more easily acted upon by other 
substances than the crystallized. Thus, hydrofluoric acid 
acts much more readily upon them. When the amorphous 
compound is boiled with solutions of potassium or sodium 
hydroxide, or of the carbonates of these metals, it dis- 
solves, forming the corresponding silicate : 

K 2 C0 3 +Si0 1 = K t SiO,+ C0 1 ; 
2KOH + Si0 2 = K 2 Si0 3 + H 2 0. 

The crystallized varieties are not dissolved in this way. 
All forms of the dioxide act upon melting hydroxides or 
carbonates of potassium or sodium, and form the corre- 
sponding silicates. 

Uses. — Plants take up silicon dioxide from the soil, and 
this being deposited in various parts of their tissues, gives 
them the necessary firmness. Straw, for example, is rich 
in silicon dioxide. Horsetail, a plant of the genus 
Equisetum, is so rich in finely-divided silicon dioxide that 
it is used for polishing. There are great natural deposits 
of finely-divided silicon dioxide known as infusorial earth. 
This consists of the remains of diatoms. And finally sili- 
con dioxide is found in the hair, in feathers, and in egg 
albumen. Silicon dioxide finds extensive application in 
the manufacture of mortar, glass, and porcelain. Ordinary 
glass, as we shall see, is a silicate of calcium and potassium 
or sodium, which is made by melting together sand and 
the carbonates of the metals mentioned. 

Silicic Acid. — There are many varieties of siliciG acid, 
all of which can, however, be referred to the normal acid. 



SILICIC ACID. 395 

Si(OII) 4 . This normal acid is contained in the gelatinous 
precipitate which is formed when silicon tetrachloride or 
tetrafluoride is decomposed by water : 



SiCl 4 + 4H 2 = Si(OH) 4 + 4HC1. 

The normal acid readily loses water and forms the acid 
of the formula OSi(OH) 2 or H 2 Si0 3 , From the latter most 
of the ordinary silicates are derived. It cannot be isolated, 
for when filtered off and exposed to the air it loses more 
water, and when heated to a sufficiently high temperature 
it is converted into silicon dioxide. 

When potassium or sodium silicate in solution is treated 
with hydrochloric acid, most of the silicic acid separates 
in the form of a gelatinous mass if the solution is concen- 
trated. If, however, the solution is dilute, a considerable 
part of the acid remains in solution. Further, if a con- 
centrated solution of the silicate of potassium or sodium is 
poured quickly into hydrochloric acid, or if the acid is 
poured quickly into the solution of the silicate, the silicic 
acid remains in solution. If, however, the solutions are 
brought together drop by drop the silicic acid separates. 
From these solutions of silicic acid ammonia or ammonium 
carbonate throws down the acid. 

A solution of pure silicic acid can be obtained by means 
of dialysis. It has been found that solutions of different 
substances pass with different degrees of ease through 
porous membranes, just as gases differ as regards the ease 
with which they pass through porous diaphragms. This 
fact concerning gases was referred to in connection with 
hydrogen. Now, while some solutions pass readily through 
parchment paper, others pass through with difficulty, and 
some do not pass through at all. A dialyser, or an 
apparatus used in dialysis, may be made by tying a piece 
of parchment paper over the mouth of a ring-formed glass 
or rubber vessel, and placing this in another shallow 
vessel. Pure water is put in the outer vessel, and the 



39^ 



COLLEGE CHEMISTRY. 



solution for dialysis in the inner one. The arrangement 
is illustrated in Fig. 76. 

In the figure act is the hoop of gutta-percha, and b is 
the parchment paper. When now the solution containing 
hydrochloric acid, sodium chloride, and silicic acid is put 
in the dialyser, the hydrochloric acid and sodium chloride 




pass readily through the membrane, while the silicic acid 
is left behind, and in the course of a few days, if the water 
in the outer vessel is renewed, the solution of silicic acid 
in the inner vessel will be found to be free from the other 
substances. This solution can be evaporated to some 
extent by boiling, but when a certain concentration is 
reached the acid separates. In a vacuum such a solution 
can be evaporated further without the formation of a 
deposit. Finally, there is left a transparent mass which 
has approximately the composition represented by the 
formula H 2 Si0 3 . The dialysed solution of silicic acid is 
coagulated by a very dilute solution of sodium or potassium 
carbonate, and by carbon dioxide itself. 

When the solutions containing silicic acid are evaporated 
to complete dryness the acid is converted into silicon 
dioxide and insoluble hydrates. This residue is called 
insoluble silicic acid. When this is treated with hydro- 
chloric acid and water it remains undissolved, and if 
filtered off and ignited it leaves a residue of silicon dioxide. 
To sum up, then: Whenever silicic acid is formed in a 
solution it is a more or less complex derivative of normal 






POLY SILICIC AND TRTSILICIC ACIDS. 397 

silicic acid, and is somewhat soluble in water, but by the 
processes just described the soluble acid is converted into 
insoluble silicon dioxide, as explained. 

Polysilicic Acids. — Silicic acid is remarkable for the 
great number of derivatives which it yields. Most of these 
bear to the normal acid relations similar to those which the 
various forms of phosphoric acid bear to normal phos- 
phoric acid. It has already been stated that salts of the 
acid H 2 Si0 3 are more common than those of the normal 
acid. Among the salts of the normal acid are zircon, 
ZrSi0 4 , and thorite, ThSi0 4 . The ordinary silicates of 
potassium and sodium are derived from the acid H 2 Si0 3 ; 
so also are wollastonite, CaSi0 3 , and enstatite, MgSi0 3 . 

Disilicic Acid is derived from ordinary silicic acid by 
loss of one molecule of water from two molecules of the 
acid: 

.OH 

OH 0Si \ 
20Si<on= >° + H 2 0. 
OSi< 
x OH 

Its composition is, therefore, H 2 Si 2 5 , which may be 
written 3 Si 2 (OH) 2 . Another form of disilicic acid is 
derived from two molecules of the normal acid by loss of 
one molecule of water: 

2Si(OH), = OSi,(OH), + H 2 0. 

The well-known mineral sepentine is apparently the 
magnesium salt of this acid. It is represented by the 
formula Mg 3 Si 2 O r 

Trisilicic Acids are derived from three molecules of the 
normal acid or the ordinary acid by loss of different num- 
bers of molecules of water. Thus, by loss of two mole- 
cules the normal acid would yield a product H 8 Si 3 O 10 . By 
loss of two molecules of water this trisilicic acid Would 
yield an acid of the formula H Si o 8 . The structure of 



39 8 COLLEGE CHEMISTRY. 

the first acid is expressed by formula I, and of the second 
by formula II, below given : 



Si )(0 H )» Sijo Sijo 

Si I(OH), Sijg (0H ), SM0 3 A1 

I. II. III. 

Orthoclase or ordinary feldspar is the aluminium-potas- 
sium salt of the second form of trisilicic acid, in which 
one atom of hydrogen is replaced by potassium, and three 
by an aluminium atom, as shown in formula III above. 

Titanium Dioxide, Ti0 2 . — As has been stated, this is 
one of the principal forms in which titanium is found in 
nature. There are three natural crystallized varieties — 
rutile, brookite, and anatase. 

Silicides are compounds of silicon with other elements, 
as, for example, with carbon. These two elements com- 
bine, forming an interesting compound, carton silicide, 
CSi, which is manufactured on the large scale and known 
in the market as carborundum. This is made by heating 
a mixture of quartz sand, coke, and common salt, or 
sodium chloride, in the electric furnace to 3500°, when 
the reaction represented below takes place : 

Si0 2 + 2C = CSi + 2CO. 

The product is in the main crystallized, the crystals 
being bluish or yellowish-green. They have the specific 
gravity 3.22 to 3.12. The silicide is said to be colorless 
when perfectly pure. It scratches ruby and chrome-steel, 
and on account of its hardness it is much prized as a 
polishing agent, being used to a considerable extent in 
place of emery. Pure carbon silicide is insoluble in nearly 
all ordinary solvents, including hydrochloric, nitric, sul- 
phuric, and hydrofluoric acids. It is, however, decom- 
posed by fusing caustic alkalies or their carbonates. 



EXPERIMENTS WITH SILICON. 399 

Many other silicides have been made, several of which 
are well-crystallized compounds. 

Family IV, Group B. 

Allied to the members of the silicon group, yet differing 
from them in some important particulars, are the three 
elements germanium, tin, and lead. Of these the first 
two are more acidic in character than the last. They 
combine with chlorine in two proportions, forming the 
chlorides GeCl 2 , SnCl 2 , PbCl 2 , GeCl 4 , SnCl 4 , PbCl 4 . 
With oxygen they unite, forming the compounds Ge0 2 , 
Sn0 2 , and Pb0 2 . Stannic oxide, Sn0 2 , and lead peroxide, 
Pb0 2 , form salts with bases, and these have the composi- 
tion represented by the general formulas M,SnO s and 
M a Pb0 3 , and are therefore analogous to the silicates and 
titanates. On the other hand, further, salts are known 
which are derived from the oxide PbO. These have the 
general formula M 2 Pb0 2 , and are to be regarded as salts 
of an acid, Pb(OH) 2 . These salts are not stable, and are 
not easily obtained. Most of the derivatives of lead are 
those in which it plays the part of a base-forming element. 
It will therefore be better to postpone its study until it is 
taken up under the general head of the base-forming 
elements. Notwithstanding, further, the marked analogy 
between some of the compounds of tin and those of the 
members of the silicon group, it appears on the whole 
advisable to treat of this element in company with lead, 
which it also resembles in many respects. 

EXPERIMENTS. 

Silicon. 

Experiment 196. — Prepare sodium fluosilicate as directed in 
the next experiment. Mix 3 parts of the dry salt with 1 part of 
sodium cut in pieces. Throw this mixture all at once into a Hes- 
sian crucible heated to bright red heat in a furnace. Add imme- 
diately 9 parts granulated zinc, and a layer of sodium chloride 
previously heated to drive off water. The crucible is then cov- 



4oo 



COLLEGE CHEMISTRY. 



ered, and the fire allowed to burn down. After cooling, the 
regulus of zinc containing the silicon is separated from the slag, 
washed with water, and treated with hydrochloric acid. The 
zinc dissolves and leaves the silicon. This is again washed with 
water, and then heated with nitric acid, and washed with water, 
when crystals of silicon, sometimes of great beauty, are obtained. 
Try the effect of heating a little of the silicon in the air. Try the 
action of acids and of alkalies upon it. 

Silicon Tetrafluoride and Fluosilicic Acid. 

Experiment 197. — Arrange an apparatus as shown in Fig. 77. 
A is a bottle of about 2 litres capacity, such as are commonly 
used for transporting acids. This is about two-thirds filled with 




"xT 7 



Fig. 77. 



alternating layers of sand and powdered fluor-spar, moistened 
with concentrated sulphuric acid. The bottle is put in the deep 
sand-bath B, and connected by means of a wide glass tube with 
the funnel G, which dips just below the surface of the water in 
the large evaporating dish D. The sand-bath is now gently 
heated, when silicon tetrafluoride passes over. Coming in con- 
tact with water, it is decomposed, silicic acid being deposited and 
fluosilicic acid passing into solution. In order to prevent clog- 
ging, the gelatinous silicic acid is from time to time removed 
from the mouth of the funnel by means of a bent glass rod. 
After the action is complete, filter the solution. Take out one- 



EXPERIMENTS WITH SILICIC ACID. 401 

quarter, and to the rest slowly add a solution of sodium carbon- 
ate until the whole just begins to show an alkaline reaction ; now 
add the other quarter of the acid, and filter. Explain all the re- 
actions. Heat a little of the dried salt in a covered platinum 
crucible. What change takes place ? What evidence have you 
that the change has taken place ? To a little of the salt in water 
add a solution of potassium hydroxide. What change takes place ? 
Dry the silicic acid formed in the first part of the experiment by 
decomposition of the silicon tetrafluoride. 



Silicic Acid. 

Experiment 198.— Boil some of the silicic acid obtained in 
the last experiment in a solution of sodium hydroxide. Treat 
some of the solution with hydrochloric acid ; with ammonium 
chloride. 

Experiment 199. — Fuse a mixture of equal weights potassium 
and sodium carbonates in a platinum crucible in the flame of the 
blast-lamp and slowly add fine sand to the molten mass. Con- 
tinue the heating until no more sand is dissolved. Pour the 
molten mass out on a stone, and when cooled break it up and treat 
it with water. 

Experiment 200.— Treat a little of the solution containing 
sodium and potassium silicates, prepared in the last experiment, 
with a little sulphuric or hydrochloric acid. A gelatinous sub- 
stance will be precipitated. This is silicic acid. Some of the 
acid remains in solution. By evaporating the solution to dryness 
and heating for a time on the water-bath, all the silicic acid is 
rendered insoluble. 



CHAPTER XXIII. 

CHEMICAL ACTION. 

Retrospective. — We have been studying the principal 
elements of four families and the compounds which they 
form with one another. No matter how simple or how 
complex the chemical changes studied were, certain funda- 
mental laws governing all cases of chemical action were 
found to hold good. These laws have been discussed, but 
it will be well to recall them here before taking up others 
which are intimately connected with them. The first 
great law of chemical change is 

I. The law of conservation of mass. 

According to this the amount of matter is not changed 
by a chemical act. 
The second law is 

II. The law of definite proportions. 

According to this, the composition of any compound is 
always the same. 
The third law is 

III. The law of multiple proportions. 

According to this, the different masses of any element 
that combine with. a fixed mass of another or others bear 
simple relations to one another. 

To account for the laws of definite and multiple propor- 
tions the Atomic Theory has been proposed. 

According to this, each element is made up of particles 
of definite weight, which are chemically indivisible, and 

402 



CHEMICAL ACTION. 4°3 

chemical action consists in union or separation of these 
particles. These hypothetical particles are called atoms. 
The elements must combine in the proportion of their 
atomic weights or of simple multiples of these, if the 
atomic theory is true. 

Further study showed that it is necessary to assume the 
existence of larger particles than the atoms, viz., the mole- 
cules. According to the theory of molecules, every chem- 
ical compound and element is made up of molecules, which 
are the smallest particles having the same general proper- 
ties as the mass. These molecules are made up of atoms 
which, in the case of compounds, are of different kinds, 
and in the case of elements, of the same kind. In the 
case of a few elements the atom appears to be identical 
with the molecule. 

From the study of gases the conclusion is reached that 
in equal volumes of all gases under standard conditions 
there is always the same number of molecules (Avogadro's 
hypothesis). This furnishes a means of determining the 
relative weights of molecules of gaseous substances; and from 
these molecular weights it is possible to draw conclusions 
in regard to the atomic weights of those elements which 
enter into the composition of the compounds thus studied. 

The formulas of chemical compounds are intended to 
be molecular formulas. They are intended to tell of what 
atoms and of how many atoms the molecules represented 
are made up. 

The method of determining molecular weights based 
upon Avogadro's hypothesis is applicable only to gaseous 
substances, or to such as can be converted into gas without 
undergoing decomposition. While many of the com- 
pounds with which we have had to deal are of this 
character, many of them are not, and in regard to the 
molecular weights of these we must be in doubt unless 
some other method applicable to liquids and solids is 
available. So, too, the atomic weights of those elements 
which enter into the composition of gaseous compounds 
can be deduced from the molecular weights, but plainly 



4°4 COLLEGE CHEMISTRY. 

those which do not enter into the composition of snch 
compounds demand some other method. For determining 
the atomic weights of such elements an excellent method 
is based upon the study of specific heats; while for the 
determination of the molecular weights of solid substances 
which can be dissolved without decomposition a method 
has been worked out that is based upon the extent to which 
the compound raises the boiling-point or lowers the freez- 
ing-point of its solution. Both these methods will be 
briefly described in this chapter. 

Next, it is found that there is a limit to the law of 
multiple proportions. While, according to this law, the 
masses of any element which unite with a given mass of 
another element bear simple relations to one another, the 
law is silent as to how many kinds of compounds are 
possible between any two elements. A careful examina- 
tion of the composition of the compounds of the elements 
shows, however, that there is a limit to the n amber of 
atoms of one element that can combine with one atom of 
another element. This limit is determined by what is 
called the valence of the elements. Observations on the 
composition of compounds have led to the hypothesis of 
the Unking of atoms — the linking taking place according 
to the laws of valence. The arrangement of the atoms in 
a molecule is, acording to this, the constitution of a com- 
pound. 

Valence, as we have seen, is not a constant property of 
the atoms. Towards oxygen the elements which we have 
thus far studied have the highest valence; towards hy- 
drogen the lowest; and, in general, towards the members 
of the chlorine group they exhibit an intermediate valence. 
The valence towards hydrogen is in most cases constant, 
while the valence towards oxygen and towards the mem- 
bers of the chlorine group varies, in some cases between 
comparatively wide limits, as between 1 and 7 in the 
chlorine group, and between 2 and 6 in the sulphur group. 
Further, the variations in the valence of an element 
generally take place from odd to odd or from even to even. 



CLASSIFICATION OF REACTIONS. 4^5 

In the case of chlorine it appears to vary from 1-3 to 5-7; 
in that of sulphur, from 2-4 to 6 ; in that of phosphorus, 
from 3 to 5. 

A comparison of the atomic weights finally led to the 
discovery that the properties of the elements are a periodic 
function of these weights. This is the periodic law of 
chemistry. This makes a systematic classification of the 
elements according to their atomic weights and their 
properties possible, and is full of suggestion as to the 
relations which the forms of matter we call elements bear 
to one another. 

Classification of Reactions of the Elements and Com- 
pounds Studied. — While there is undoubtedly something 
confusing in the number of the compounds and their 
reactions which we have been studying, still, when these 
are interpreted in the light of the atomic theory, of the 
law of valence, and of the periodic law, the study is much 
simplified, and those things which seem to have little or 
no connection are found to form parts of a general system. 
In studying chemistry, one of the first things to be done is 
to learn how elements and compounds act upon one 
another, and what products are formed. The question of 
composition is one of the first that presents itself, and this 
must be studied before other questions can be intelligently 
discussed. What, then, are the most prominent facts 
which we have learned in studying the elements and com- 
pounds which have thus far been taken up ? 

Kinds of Chemical Reactions. — All chemical reactions 
may be classified under three heads: 

(1) Those which consist in direct combination; 

(2) Those which consist in direct decomposition; and 

(3) Those which involve the interaction of two or more 
elements or compounds and the formation of two or more 
compounds. This is known as double decomposition or 
metathesis. 

Direct Combination. — We have had to deal with a num- 
ber of examples of each of these kinds of reactions. As 



4<>6 COLLEGE CHEMISTRY. 

examples of the first kind already studied the following 
may be mentioned : 

The combination of hydrogen and chlorine to form 
hydrochloric acid; the formation of ammonium chloride 
from ammonia and hydrochloric acid; the formation of 
calcium hydroxide from calcium oxide and water; the 
formation of nitrogen peroxide from nitric oxide and 
oxygen; and the formation of carbon disulphide from 
carbon and sulphur. 

As regards the combination of hydrogen and chlorine, 
it should be remarked that this act is the same in principle 
as that of metathesis. Strictly speaking, it is not a case 
of direct combination, as we understand it. For, as we 
have seen, according to the molecular theory, free chlorine 
and free hydrogen consist of molecules which are made up 
of two atoms each. Therefore, when these elements are 
brought together the molecules are first decomposed into 
atoms before the act of union can take place. The two 
acts are represented by the two equations following : 

C1 2 +H 2 = C1+C1 + H + H; 
CI + CI + H + H = 2HC1. 

In the case also of the union of hydrochloric acid and 
ammonia it appears probable that a serious disarrangement 
of the constituent atoms is necessary in order that the act 
of combination may take place. According to the am- 
monium theory, ammonium chloride is represented by the 
PH 

H 
formula N \ H, which means that the atom of chlorine 

H 

CI 

and four atoms of hydrogen are in combination with the 
atom of nitrogen. But in order that a compound of this 
constitution may be formed from ammonia and hydro- 
chloric acid, it is necessary that the molecule of hydro- 
chloric acid should be broken down into its constituent 
atoms. So that this case of apparent direct combination 



DIRECT DECOMPOSITION. 4° 7 

is, as far as we can judge, in reality more complicated than 
it appears, and should be represented by the two equations : 

NH 3 + HC1 = NH 3 + H + CI; 
NH 3 + H + CI = NH 4 C1. 

All other cases of apparent direct combination are prob- 
ably of the same character, so that it is doubtful whether 
a single case of simple direct combination is known. 

Direct Decomposition. — As examples of direct decom- 
position the following cases may be cited : 

The decomposition of mercuric oxide by means of heat 
into mercury and oxygen; that of ammonium chloride into 
ammonia and hydrochloric acid by heat ; that of potassium 
nitrate into potassium nitrite and oxygen by heat ; that of 
phosphorus pentachloride into the trichloride and chlorine 
by heat; that of ammonia into hydrogen and nitrogen by 
continued action of electric sparks; and that of nitrogen 
iodide by contact with a solid substance. 

On close examination of each of the above cases, which 
are fairly typical and as simple as any that could be 
chosen, it will be seen that no one of them is merely a case 
of decomposition; for even though we must assume that 
the first result in each case is the setting free of the atoms 
of one or two elements, we must also assume that these 
atoms unite again to form other molecules either of ele- 
ments or compounds. Thus, when mercuric oxide is 
decomposed we get mercury and oxygen. As far as can 
be determined, the mercury atoms do not unite with each 
other, but the oxygen atoms do, so that the total action 
involves decomposition and afterwards combination as 
represented in the equations 

2HgO = Hg + Hg + + 0; 

Hg + Hg + + = Hg + Hg + 2 . 

In the case of the pentachloride of phosphorus, it is prob- 
able that the two atoms of chlorine are first given off from 
each molecule of the chloride, leaving a molecule of the 



408 COLLEGE CHEMISTRY. 

trichloride, but the atoms of chlorine afterwards unite to 
form molecules as represented thus : 

PC1 8 = PC1 8 +C1 + C1; 
PC1 3 + C1+C1 = PC1 3+ C1,. 

Similar statements hold good for all other cases of direct 
decomposition. 

Metathesis. — This is the most common kind of chem- 
ical action, and indeed from what has been said in regard 
to direct combination and direct decomposition it will be 
seen that there is no essential difference between them and 
metathesis. Most of the reactions with which we have 
had to deal are examples of double decomposition or 
metathesis, as: The formation of salts by the action of 
bases upon acids; the formation of the sulphides of arsenic, 
antimony, and bismuth by the action of hydrogen sulphide 
upon solutions of compounds of these elements; the 
setting free of hydrochloric and nitric acids by the action 
of sulphuric acid upon chlorides and nitrates; of carbon 
dioxide and oxides of nitrogen by the action of acids upon 
carbonates and nitrites; and of ammonia by treating am- 
monium salts with lime. As simple an example of this 
kind of action as can be given is that of the formation of 
hydrogen and potassium chloride from potassium and 
hydrochloric acid gas , The molecular weight of potassium 
is not positively known, but, assuming its molecule to be 
made up of two atoms, the action must be represented in 
this way: 

K 2 + 2HC1 = 2KC1 + H 2 . 

The next stage of complication is exhibited in the 
reaction following: 

KI + HC1 = KC1 + HI. 

Examples similar to the latter, but somewhat more com- 
plicated, are these: 

2KOH + H 2 S0 4 = K 2 S0 4 + H 2 0; 
CaCl 2 + H 2 S0 4 = CaS0 4 + 2H01. 



THE CAUSE OF CHEMICAL REACTIONS. 409 

The Cause of Chemical Reactions. — The prime cause of 
chemical reactions is something which we think of as an 
attractive force exerted in different degrees between the 
different elements. When any elements or compounds 
are brought together under certain conditions the tendency 
is always towards the formation of the most stable com- 
pounds of those elements which can be formed under the 
given conditions. Thus, potassium sulphate and water 
are more stable forms of combination of the elements 
hydrogen and oxygen, and potassium, sulphur and oxygen, 
than sulphuric acid and potassium hydroxide are under 
the conditions under which the action takes place. So also 
the system composed of potassium chloride and hydriodic 
acid is more stable than that composed of potassium iodide 
and hydrochloric acid under the conditions of the action. 
Why the one system is more stable than the other we do 
not know, for we do not know what relations exist between 
the atoms in the molecules. It is convenient to think of 
that which causes the atoms to unite to form compounds 
as an attractive force. It is evident that this force is more 
strongly exerted between some elements than between 
others. That between chlorine and hydrogen is, for ex- 
ample, much stronger than that between chlorine and 
nitrogen or oxygen. Owing, however, to the complicated 
character of most chemical reactions, it is extremely diffi- 
cult to make measurements of this force. 

An Ideal Chemical Reaction. — In every case in which 
two compounds act upon each other to form two new ones, 
several forces must be at work, as we have seen. Suppose, 
for example, AB and CD act upon each other in the 
gaseous condition to form two compounds BC and AD, 
also both gaseous. The normal course of such a reaction 
would lead to the formation not only of the two compounds 
BC and AD, but AB and CD would also be present in the 
resulting system. For A has the power to combine with 
B as well as with D, and C has the power to combine with 
D as well as with B. In the system we should have 
operating the tendency of A to combine with B, and that 



410 COLLEGE CHEMISTRY. 

of A to combine with D\ that of to combine with D, 
and that of C to combine with B. As these forces operate 
simultaneously, equilibrium is established when certain 
quantities of the four possible compounds are formed, the 
quantities depending in the first instance upon the relative 
strengths of the forces at work. 

It is difficult to give an illustration of this kind of rela- 
tion, because the study of chemical changes between gases 
that give rise to the formation of gaseous products is diffi- 
cult, and the study of changes in solution is complicated 
by the action of the solvent and, if the solvent is water 
and the substances are acids, bases, or salts, this action 
plays a controlling part. 

Influence of Mass. — When hydrogen sulphide is passed 
into a solution of cadmium chloride, cadmium sulphide is 
precipitated, the reaction being represented by the equa- 
tion 

(1) CdCl 2 + H 2 S = CdS + 2HC1. 

On the other hand, when hydrochloric acid is added to 
cadmium sulphide the latter dissolves and the action is 
represented by the equation 

(2) CdS + 2HC1 = CdCl 2 + H 2 S. 

This action is plainly the reverse of that which takes place 
when hydrogen sulphide acts upon cadmium chloride in 
solution. The difference between the two cases is due to 
the relative masses of hydrogen sulphide and of hydro- 
chloric acid that are brought into action. If the hydro- 
chloric acid is very dilute and the hydrogen sulphide is 
present in excess, reaction (1) takes place and is complete. 
If the conditions are reversed, then reaction (2) takes 
place. If considerable hydrochloric acid is added to a 
solution of cadmium chloride, and hydrogen sulphide 
passed into the solution, all of the cadmium will not be 
precipitated as sulphide, and the amount precipitated will 
depend upon the amount of hydrochloric acid present. 



INFLUENCE OF MASS. 4^ 

A reaction that can take place in two directions or that is 
reversible is commonly represented thus : 

CdCl, + H 2 S J£ CdS + 2HC1. 

Equilibrium is established in such cases under definite 
conditions, and the law governing the equilibrium has 
been discovered. 

Guldberg and Waage were the first ones to formulate 
the law of mass action. The law is as follows: 

When tivo substances act upon each other, the action is 
proportional to the active mass of the substances talcing part 
in the change. 

By active mass is meant the molecular concentration, 
or, in other words, the number of gram-molecules per 
litre, a gram-molecule being the number of grams corre- 
sponding to the molecular weight of the substance. A 
gram-molecule of sodium chloride, for example, is 58.5 
grams, etc. 

The following quotation from the memoir by Guldberg 
and Waage will make the subject somewhat clearer : 

" When two substances A and B are transformed into 
two new substances A' and B', the chemical force with 
which A and B act upon each other is measured by the 
quantity of the new substances formed in unit time. 

"The quantity of a substance in unit volume of the 
compound in which the chemical change takes place we 
call the active mass of the substance. 

" The chemical force with which two substances A and 
B act upon each other is equal to the product of their 
active masses, multiplied by the coefficients of affinity. 

"By coefficient of affinity a coefficient is understood 
which is dependent upon the chemical nature of the two 
substances and upon the temperature. If the active 
masses of A and B are represented by^? and q, and the 
coefficient of affinity by k, then the force acting between 
A and B is expressed by Jcpq. 

" When in a chemical process A and B are transformed 



4** COLLEGE CHEMISTRY. 

into A' and B', and A' and B' can at the same time be 
transformed into A and B, equilibrium will be established 
when the force acting between A and B is equal to the 
force acting between A' and B'. 

" If the active masses of A' and B' are represented by 
p' and q' and their coefficient of affinity by k', the chemical 
force acting between A' and B' is expressed by k'p'q'. 

"The condition of equilibrium is therefore expressed 
by the equation hpq — k'p'q'." 

Reactions may be Complete if one of the Products 
Formed is Insoluble or Volatile. — When two substances 
which by interaction can form an insoluble product are 
brought together, the reaction generally takes place and 
is complete. When the substances are brought together 
we may imagine that, owing to interaction, a small quan- 
tity of the insoluble compound is formed at once. If this 
product were soluble, the action would stop before it is 
complete, because this new product would itself exert its 
action upon the system. Being insoluble, however, it is 
removed from the sphere of action, and the same reaction 
which caused the formation of the first particles of it can 
now be repeated, and so on, until the reaction is complete. 
This is illustrated in the action of sulphuric acid upon 
barium chloride in solution. The two substances react as 
represented in this equation : 



BaCl, 



+ H 2 SO, + Aq = BaS0 4 + HC1 + Aq. 



The symbol Aq is simply intended to indicate that the 
reaction takes place in solution. If barium sulphate were 
soluble, all four substances — barium chloride, sulphuric 
acid, barium sulphate, and hydrochloric acid — would be 
present in the solution after the establishment of equili- 
brium. But, being insoluble, it is removed, and new 
quantities are formed as long as the substances necessary 
for its formation are present in the solution; that is, until 
either all the barium chloride is decomposed or all the sul- 
phuric acid is removed. Reactions involving the formation 



COMPLETE REACTIONS. 4*3 

of insoluble compounds or precipitates are among the most 
common with which we have to deal, particularly in the 
various operations of analytical chemistry. 

Again, when two substances which can form a volatile 
product are brought together the reaction generally takes 
place and is complete. The reason why a reaction of this 
kind is complete is the same as that given in the case of 
the formation of an insoluble compound. Each successive 
portion of the volatile product formed is removed, and the 
reaction which gave rise to it proceeds as long as the 
necessary substances are present. This kind of action has 
been repeatedly illustrated. It is that, for example, which 
is seen in the liberation of hydrochloric acid from a chlo- 
ride by the action of sulphuric acid ; of carbon dioxide by 
the action of an acid upon a carbonate; and of ammonia 
by the action of lime upon ammonium chloride. 

An interesting example of the combined influence of 
mass and the volatility of the product is seen in the action 
of heated iron upon an excess of steam, and of the oxide 
of iron upon an excess of hydrogen. When steam is 
passed over heated iron, action takes place thus : 

4H 2 + 3Fe = Fe 3 4 + 4H 2 . 

Hydrogen is liberated and the oxide of iron formed. 
When, however, hydrogen is passed over heated oxide of 
iron the reverse reaction takes place : 

Fe 3 4 + 4H 2 = 3Fe + 4H 2 0. 

Owing to the excess of steam always present in the first 
reaction, hydrogen is constantly formed and constantly 
being removed. Undoubtedly the hydrogen formed acts 
to some extent upon the oxide, but the other reaction 
always takes place to a greater extent. The opposite is 
true when the oxide is heated in an excess of hydrogen. 
The principal reaction which takes place in this case is 
that of the hydrogen upon the oxide of iron, and the steam 
is carried out of the field almost as soon as formed, so that 
the reduction of the oxide of iron continues. 



414 COLLEGE CHEMISTRY. 

Dissociation. — It has already been stated (p. 87) that 
when water-vapor is heated to a sufficiently high tempera- 
ture it is decomposed into the two gases hydrogen and 
oxygen. This decomposition does not take place suddenly 
when a certain temperature is reached, but it begins at a 
comparatively low temperature and increases as the tem- 
perature is raised. For any given temperature and 
pressure the amount of decomposition is definite. On 
cooling a mixture of water-vapor and hydrogen and 
oxygen, the two elements combine to a greater and greater 
extent as the temperature is lowered. A condition of 
equilibrium is established under each new set of condi- 
tions. The gradual decomposition of a substance into its 
constituents which reunite when the temperature is 
lowered is called dissociation. All cases of dissociation 
are examples of reversible reactions, and they also illustrate 
the phenomena of equilibrium. Other examples besides 
that mentioned are the change of nitrogen peroxide from 
the form N 2 4 to that of N0 2 . This may be represented 
thus : 

NA ~ N0 2 + N0 2 . 

The dissociation of hydriodic acid into hydrogen and 
iodine may also be mentioned : 

8ffl£H,+T, 

In all of these cases the substances taking part are gaseous. 
When, however, calcium carbonate breaks down into 
lime and carbon dioxide two of the three substances are 
solids and one is a gas : 

CaC0 3 ^CaO + C0 2 . 

Again, the dissociation of ammonium chloride into am- 
monia and hydrochloric acid involves the interaction of 
three substances, two of which are gases and one a solid : 

NH 4 C1 J£ NH 3 + HOI, 



DISSOCIATION. 415 

It would lead too far to discuss these reactions m detail 
here. 

The explanation of the phenomena of dissociation of 
gases is found in the kinetic theory of gases. According 
to this theory, the molecules of a gas at a given tempera- 
ture are. moving with different velocities, though the 
average velocity of all the molecules is the same at the 
same temperature. Now, it is highly probable that the 
motion of the atoms within the molecules partakes of that 
of the molecules themselves, so that the motion of the 
atoms in the molecules with the greatest velocity is prob- 
ably the greatest, and, in these, decomposition will take 
place first. When a compound gas is heated, we can easily 
conceive that even at a comparatively low temperature the 
motion of some of the molecules will be sufficient to cause 
their decomposition, and, as the average motion of all the 
molecules is constant for a given temperature, the amount 
of decomposition will be constant for that temperature. 
As the molecules are, however, moving in every direction 
and constantly colliding, a molecule which is decomposed 
at one instant may be re-formed at the next, and one that 
is not decomposed may acquire motion enough to cause its 
decomposition. Though, as is believed, these changes are 
constantly taking place at every temperature, still, as has 
been said, the number of molecules which will be decom- 
posed in a given mass at a given temperature and pressure 
will always be the same. The higher the temperature, 
then, the greater the number of molecules in the condi- 
tions which cause decomposition, and the smaller the 
number of those in the conditions favorable to formation. 
At each temperature and pressure an equilibrium is estab- 
lished, the number of molecules decomposed being equal 
to the number formed. It is obvious that, if one of the 
products of decomposition is removed, the conditions are 
entirely changed. Then the possibility of recombination 
will not exist, and total decomposition can be effected at 
a lower temperature than that required for total decom- 
position in the process of dissociation proper. 



416 COLLEGE CHEMISTRY. 

Electrolysis. — Some chemical compounds in solution in 
water conduct electricity, and at the same time they 
undergo decomposition. Thus, hydrochloric acid in solu- 
tion in water conducts electricity, and its constituents 
hydrogen and chlorine are obtained, the hydrogen appear- 
ing at the negative and the chlorine at the positive pole. 
Compounds that act in this way are called electrolytes. 
When a current of electricity acts upon solutions of differ- 
ent salts, equivalent quantities of the metals are deposited 
by the same current in the same time. This is Faraday's 
Law. Thus if the same current were passed simul- 
taneously through solutions of silver nitrate, AgN0 3 , 
mercuric nitrate, Hg(N0 3 ) 2 , cupric sulphate, CuS0 4 , and 
ferric chloride, FeCl 3 , it would be found that for every 
107.93 parts by weight of silver deposited there would be 
100.15 parts by weight of mercury deposited, 31.8 of 
copper, and 18.67 of iron. These are equivalent quantities 
of these metals — quantities that take the place of one part 
by weight of hydrogen — and are to be distinguished from 
atomic quantities. Those elements which appear at the 
negative pole, or cathode, are called electro -positive, and 
those which appear at the positive pole, or anode, are 
called electro-negative. Those elements which we call 
acid-forming are electro-negative, while hydrogen and the 
base-forming elements are electro-positive. The ions that 
tend towards the positive pole, or anode (the acid-form- 
ing, or electro-negative elements, etc.), are the anions, 
while those which tend towards the negative pole, or 
cathode (base-forming, or electro-positive elements, etc.), 
are the cations. The electrolysis of chemical compounds is 
not generally a simple decomposition into two constituents. 
Thus, when copper sulphate, CuS0 4 , is decomposed, the 
copper is deposited at the negative pole ; but no such com- 
pound as S0 4 appears at the positive pole. This, if formed, 
forms sulphuric acid and gives off oxygen. Both oxygen 
and sulphuric acid as a matter of fact appear at the positive 
pole. The changes involved may be represented thus : 
CuS0 4 = Cu + S0 4 ; S0 4 + H,0 = H,S0 4 + 0, 



ELECTROLYTIC DISSOCIATION 417 

Electrolytic Dissociation. — It has been known for a long 
time that a very weak electric current acting upon a solu- 
tion of an electrolyte is sufficient to cause the ions to 
appear at the poles. This fact is inexplicable if it is 
assumed that the current is the cause of the decomposition 
of the electrolyte. This, and some other facts which will 
be referred to farther on, make it probable that electrolytes 
are at least to some extent decomposed into their constit- 
uent ions when they are dit solved in water; that these ions 
charged with electricity transfer their charges in the solu- 
tion and thus conduct the current; and that when an ion 
charged with negative electricity reaches the positive pole 
its electricity is discharged, and the ion then ceases to be 
an ion and becomes an element in the free state or some 
compound which appears either as such or in the form of 
other products. According to this conception, the act of 
solution of an electrolyte, in water at least, involves partial 
breaking down or dissociation of the compound into its 
ions. The extent of this breaking down is determined 
primarily by the concentration of the solution — the greater 
the dilution the greater the dissociation. At infinite dilu- 
tion there is complete dissociation. A water solution of 
hydrochloric acid containing 36.18 grams of the acid in 
1000 litres has been shown to be completely dissociated, 
or it is to be' regarded as containing hydrogen ions and 
chlorine ions. These and all other ions are carefully to 
be distinguished from the atoms or definite compounds. 
An ion always carries with it a certain charge of elec- 
tricity \ When this is discharged the ion becomes either 
an element or a compound in the free state. When a 
solution of one electrolyte acts upon a solution of another 
the reaction observed is probably due to the interaction of 
the ions, and it is further probable that, so far as the 
compounds are present in the undissociated condition, 
they do not act upon each other. If this view- he correct 
the reactions most familiar to us are reactions of ions, and 
not of elements or compounds. When, for example, an 
acid acts upon a base in solution it appears that, so far as 



41 8 COLLEGE CHEMISTRY. 

they react, they are in dissociated condition. Thus hydro- 
chloric acid and sodium hydroxide are to be regarded as 
acting as represented in the following equation : 

H -f CI 4- Na + OH = Na + CI + H 2 0. 

The act consists in the union of the hydroxyl ion of the 
base with the hydrogen ion of the acid to form water, the 
sodium and chlorine ions remaining as ions as in a dilute 
solution of sodium chloride. In the case of nitric acid 
and potassium hydroxide the following equation represents 
the reaction at infinite dilution: 

H + NO, + i + OH = K + N0 3 + H 2 0. 

And so also whenever the act of neutralization takes place 
there is simply a union of hydrogen ions with hydroxyl 
ions to form water. 

This conception finds strong confirmation in the fact 
that the heat evolved in neutralizing equivalent quantities 
of all acids at infinite dilution is always the same — a fact 
difficult to explain if it is assumed that in the act of 
neutralization a salt is formed in the solution. 

In the following table the heats of neutralization of a 
few acids and bases are oiven : 



Acid and Base. 




Heat of 
Neutral. 


Hydrochloric acid and sodium hyd 


roxide . . 


. . 13,700 


a it 


a 


lithium 


a 


. . 13,700 


tt a 


tt 


potassium 


a 


. . 13,700 


it ft 


a 


barium 


a 


. . 13,800 


tt it 


tt 


calcium 


tt 


.. 13,900 


Hydrobromic " 


i i 


sodium 


it 


. . 13,700 


Nitric 


a 


a 


a 


. . 13,700 


Iodic " 


a 


a 


a 


. . 13,700 



As will be seen, the heat of neutralization is the same 
no matter what the base or what the acid may be, and as 
has been pointed out this fact is easily understood, if the 



ELECTROLYTIC DISSOCIATION. 4*9 

net of neutralization consists in the union of a hydroxyl 
ion with a hydrogen ion to form water. 

If the strength of an acid is determined by the extent 
to which it is dissociated, then of coarse those acids that 
are most readily dissociated are the strongest. By every 
method available hydrochloric and nitric acids are found 
to be the most readily dissociated and they are the 
strongest acids. The electrical method for the determina- 
tion of the strength of acids is based upon this theory. 
The applicability of this theory to the explanation of the 
most common reactions that take place in solution is at 
present attracting the attention of chemists. Those reac- 
tions which are made use of for the purpose of detecting 
the presence of the various elements appear in fact, as has 
already been stated, to be reactions of the ions, and when 
these ions are not present the reactions are not observed. 
For example, when silver nitrate in solution is added to 
sodium chloride in solution a precipitate of silver chloride 
is formed, the reaction taking place as represented in this 
equation : 

AgN0 3 + NaCl = AgCl + NaNO,. 

This reaction seems, however, to be due to the fact that 
ions of silver and of chlorine are present. These coming 
together form the insoluble molecules silver chloride which 
is then precipitated: 

Ag + N0 3 + m + CI = AgCl + Na + N0 3 . 

There are many compounds that contain chlorine and 
yet do not give a precipitate of silver chloride when treated 
with silver nitrate. It is believed that in these cases the 
compound is not ionised by the solvent in such a way 
as no to yield ions of chlorine, and that therefore there 
are chlorine ions present. An example of a compound 
that contains chlorine and yet does not give the reac- 
tions of this element is potassium chlorate, KC10 3 . A 
solution of this salt does not give a precipitate with silver 



4 2 ° COLLEGE CHEMISTRY. 

nitrate. It is probable that the reason of this is that the 
salt gives the ions K and C10 3 , the latter acting quite 
differently from the ion CI, as we should naturally expect. 

Definition of Acids and Bases in Terms of the Theory 
of Electrolytic Dissociation. — According to what has just 
been said it appears that when any acid is dissolved in 
water it is dissociated into hydrogen ions and something 
else. This power to yield hydrogen ions is characteristic 
of acids. On the other hand, the most marked examples 
of bases, such as sodium and potassium hydroxides, when 
dissolved in water give hydroxyl ions, and this conduct is 
characteristic of bases. 

Raoult's Methods for the Determination of Molecular 
Weights. — One great difficulty encountered in the study 
of chemical compounds is the determination of the molec- 
ular weights of those which are not gases or cannot be 
converted into vapor by heat. From some studies on the 
freezing-points of solutions, it appears that quantities of 
compounds proportional to their molecular weights cause 
the same lowering of the freezing-points, provided the 
solvent does not act chemically upon the compound. This 
fact makes it possible to determine the molecular weights 
of substances which cannot be converted into vapor, bat 
which can be dissolved. The application of the method 
is simple. Suppose water to be the solvent used. We 
know that this liquid solidifies or freezes at 0°. Now, it 
is found that by dissolving a certain quantity of some sub- 
stance in a certain quantity of water the freezing-point is 
lowered say .5°. Further, the quantities of other sub- 
stances which are necessary to lower the freezing-point of 
the same quantity of water to the same extent can be 
determined. These quantities are proportional to the 
molecular weights according to the law of Eaoult. If, 
therefore, among the substances studied there is one the 
molecular weight of which can be determined by the 
method of Avogadro, it is possible to determine x he molec- 
ular weights of all of them by the method of Raoult, as 
will readily be seen, 



DISSOCIATION OF A DISSOLVED SUBSTANCE. 421 

So, also, it has been shown that quantities of compounds 
proportional to their molecular weights cause the same 
raising of the boiling-points, provided the solvent does not 
act chemically upon the compounds or cause them to break 
down into their ions. Convenient methods have been 
devised for the determination of molecular weights of dis- 
solved substances, the methods being based upon observa- 
tions on the boiling-points and freezing-points. It should 
be noted that, when the molecular weight of a substance 
in solution has been determined, it does not follow that 
the substance has the same molecular weight when in the 
solid condition. This is a matter in regard to which 
we have little knowledge. It is quite possible that the 
molecules of solid substances may be made up of large 
aggregates of the simple molecules that probably exist 
in solutions or in vapors. There is, however, no method 
at present known that makes a determination of the com- 
plexity of these molecules, or molecular aggregates, 
possible. 

Determination of the Extent of Dissociation of a Dis- 
solved Substance. — The effect upon the boiling-point or 
freezing-point of a solution caused by the presence of a 
dissolved substance is proportional to the number of mole- 
cules in the solution or the number of individual particles, 
whether these are undecomposed molecules or the ions 
formed as a result of dissociation. Any substance that is 
dissociated in solution will give abnormal results if the 
attempt is made to determine its molecular weight by 
observations on the boiling-point or the freezing-point of 
its solutions. This method is therefore not applicable to 
solutions of electrolytes. On the other hand, the study 
of such solutions has shown that there is an increased 
lowering of the freezing-point for the same weight of 
solvent as the dilution becomes greater, a fact that points 
clearly to the conclusion that as the solution is diluted 
there is greater and greater dissociation, and advantage 
can be taken of this fact for the purpose of determining 



422 COLLEGE CHEMISTRY. 

the extent to which dissociation has taken place in a solu- 
tion of an electrolyte in water. 

In general only organic compounds come within range 
of the methods of Raonlt. These methods are now exten- 
sively used in the study of the compounds of carbon, 
simple forms of apparatus having been devised for this 
purpose. The recognition of the fact that electrolytes do 
not obey the simple law that holds good in the case of 
non-electrolytes led Arrhenius to the idea that the former 
are dissociated in solution— an idea which has proved of 
great service to the science, and is likely to revolutionize 
the views of chemists in regard to the action of chemical 
substances upon each other in solution.. 

Osmotic Pressure. — During the last century it was 
observed that, when a vessel is filled with alcohol and the 
vessel tightly covered by tying a bladder over its mouth, 
and the whole then immersed in water, the bladder is 
stretched outward, showing that the liquid from without 
has found its way into the vessel through the bladder. 
For many years investigations on phenomena of this kind 
were carried on, but no results of a general character of 
special importance from the chemical point of view were 
reached until recently. In 1877 PfefTer, the botanist, 
gave an account of a large number of experiments, and 
laid the foundation of our present knowledge of the laws 
of osmotic phenomena. The simplest phenomenon of this 
kind is seen when a wide glass tube, tightly closed with a 
piece of bladder at the lower end, is partly filled with 
alcohol and then placed in a vessel of water, so that the 
level of the liquids inside and outside the tube is the same. 
If the tube is fixed in this position, the liquid will soon be 
found to rise in the tube. There is pressure from without 
inward. This is called osmotic pressure. Pfeffer's inves- 
tigations had to do with the measurement of the pressure 
exerted by different substances under different conditions 
of temperature and dilution. Instead of using a bladder, 
which is not capable of much resistance, he used mem- 
branes made by precipitating copper ferrocyanide in the 



SPECIFIC HEAT AND ATOMIC WEIGHTS. 423 

pores of clay cells. Somewhat later Yan't Hoff showed 
that the results obtained by Pfeffer lead to the following 
remarkable laws governing osmotic pressure: 

1. The pressure is proportional to the concentration, or 
it is inversely proportional to the volume in which a definite 
quantity of the dissolved substance is contained. 

2. The pressure increases for constant volume, propor- 
tionally to the absolute temperature. 

3. Quantities of dissolved substances which are in the 
ratio of the molecular weights of these substances exert equal 
pressures at the same temperatures. 

These laws are analogous to the well-known laws of gas- 
pressure, the third being plainly analogous to the law of 
Avogadro. Another form of the third law, together with 
an extension of it, is this : 

Dissolved substances exert the same pressure, in the form 
of osmotic pressure, as they would exert ivere they gasified, 
at the same temperature, without change of volume. 

Notwithstanding the simplicity of this law, no practical 
method for determining molecular weights based upon it 
has yet been devised. The difficulties are, however, of an 
experimental nature. 

Relations between Specific Heat and Atomic Weights. 
— The fact that there is a method for the determination 
of atomic weights founded upon the relations existing 
between these weights, and the specific heat of the ele- 
ments, has been mentioned. It has been found that, when 
equal weights of different elements are exposed to exactly 
the same source of heat, they require different lengths of 
time to become heated to the same temperature. Given 
the same heating power, it requires 32 times as long to 
raise the temperature of a pound of water 10, 20, or 30 
degrees as it does to raise the temperature of a pound of 
mercury the same number of degrees; or it takes 32 times 
as much heat to raise a pound of water 10, 20, or 30 
degrees as it does to raise a pound of mercury the same 
number of degrees. Starting at the same temperature, the 
quantity of heat required to raise the temperature of a 



424 COLLEGE CHEMISTRY. 

certain weight of a substance one degree, as compared with 
the quantity of heat required to raise the temperature of 
the same weight of water one degree, is called the specific 
heat of the substance. Thus, from what was said above, 
the specific heat of mercury is £%, or, in decimals, 0.03125. 
In a similar way it can be shown that the specific heat of 
gold is 0.03244; of zinc, 0.0955; of silver, 0.057; of 
coj)per, 0.0952. 

Now, when solid elements are examined with reference 
to their specific heats, a very simple relation is found to 
exist between the numbers expressing the specific heats 
and the atomic weights. This relation will be made clear 
by a consideration of a few cases: 

Element. Specific Heat. Atomic Weight. 

Silver 0.0570 107.93 

Zinc 0.0955 65.4 

Cadmium 0.0567 112.4 

Copper 0.0952 63.6 

Tin 0.0562 118.5 

An examination of this table will show that the atomic 
weights are inversely proportional to the specific heats. 
We have 

: 0.0955 : 0.0570; 
: 0.0952 : 0.0567; 
: 0.0562 : 0.0570; etc. 



These proportions are only approximately correct; but 
it must be remembered that the means for the determina- 
tion of atomic weights and specific heats are not perfect, 
and in both sets of figures there are undoubtedly small 
errors. Hence such slight variations from absolute agree- 
ment in these proportions should occasion no surprise. 
The agreement is sufficient to indicate a close connec- 
tion between the two sets of figures. This connection 
may be stated in another way : The product of the atomic 



107.93 


65.4 : 


112.4 


63.6 : 


107.93 


118.5 : 



SPECIFIC HEAT AND ATOMIC WEIGHTS. 4*5 

weight by the specific licat is a constant. Thus, in the 
above cases: 

107.93 x 0.057 = 6.16; 
65.4 X 0.0955 = 6.24; 

112.4 x 0.0567 = 6.37; 
63.6 X 0.0952 = 6.05; 

118.5 X 0.0562 = 6.66. 

From the above it appears that the quantity of heat 
necessary to raise masses of the elements proportional to 
their atomic weights the same number of degrees is the 
same in all cases. Suppose two elements to have the 
atomic •weights 2 and 4. Their specific heats would be to 
each other as 2 to 1. That is to say, it would require 
twice as much heat to raise the temperature of a given 
mass of the element with the atomic weight 2 a certain 
number of degrees, as it would require to raise the tem- 
perature of the same mass of the element with the atomic 
weight 4 the same number of degrees. But to raise the 
temperature of masses of these two elements proportional 
to their atomic weights would require the same quantity 
of heat. This fact may be stated thns : 

The atoms of all elements have the same capacity for heat. 
This is only another way of stating that, to raise the tem- 
perature of an atom one degree, the same quantity of heat 
is always necessary. 

Now, if we assume that the constant obtained by mul- 
tiplying the specific heats by the atomic weights is 6.4, 
which is about the average of the diiferent values found, 
then it is plain that, if we divide this number by the 
specific heat of an element, we shall obtain a number which 
is very near the atomic weight. If we call the atomic 
weight A, and the specific heat H, the following equation 
expresses the relation : 

A- H . 

If this law is without exceptions, it is plain that, in order 



426 COLLEGE CHEMISTRY. 

to determine the atomic weight of an element, it is only 
necessary to determine its specific heat, and divide this 
into 6.4. The result will be very nearly the atomic 
weight. Knowing thus very nearly what the atomic weight 
is, it is a comparatively simple matter to determine it with 
great accuracy by means of chemical analysis. Unfor- 
tunately there are some marked exceptions to the law. 

Exceptions to the Law of Specific Heats. — The elements 
glucinum, carbon, boron, and silicon form exceptions to 
the law of specific heats as this law has been stated above. 
At ordinary temperatu res they do not follow the law. As 
the temperature is raised, however, the specific heat of 
these elements changes markedly, until finally,* in the 
cases of carbon and silicon, a point is reached beyond 
which there is no marked change. Thus, at 600° the 
specific heat of diamond is 0.441, and at 985° it is 0.449. 
That of silicon is 0.201 at 185°, and 0.203 at 332°. At 
these temperatures the elements obey the law. From 
elaborate studies which have been made on this subject, it 
appears that the law should be modified to read as follows : 

The specific heats of the elements vary with the tem- 
perature; but for every element there is a temperature, 
T, above which variations are very slight. The product 
of the atomic weight by the constant value of the specific 
heat is nearly a constant, lying between 5.5 and 6.5. 

Notwithstanding the irregularities referred to, the law 
of specific heats, commonly called, from the discoverers, 
the law of Dulong and Petit, is of great value in the 
determination of atomic weights. 



CHAPTER XXIV. 

BASE-FORMING ELEMENTS.— GENERAL CONSIDERA- 
TIONS. 

Introductory. — The elements thus far studied belong 
for the most part to the class of acid-forming elements, or 
those whose compounds with oxygen and hydrogen have 
acid properties. All the members of Family VII, Group 
B, are acid -forming, while the single member of Group A 
of the same family is both acid-forming and base-forming. 
All the members of Family VI, Group B, are acid-forming, 
while the members of Group A of this family are both 
acid-forming and base-forming. In Family V, Group B, 
a gradation of properties is observed, the group beginning 
with strongly marked acid-forming elements and end- 
ing with an element, bismuth, which is more basic than 
acid in character. The elements of Group A, Family 
V, are both acid-forming and base-forming, but they have 
not as sharply marked characteristics as the elements of 
Families VI and VII. Passing now to Family IV, it was 
found that the two most important members, carbon and 
silicon, belong to Group A. These two elements always 
act as acid-formers. A gradation of properties is observed 
in passing from silicon to thorium. The members of 
Group B of this family have the properties of the base- 
forming elements much more strongly marked than those 
of the acid-formers. There are still four families to be 
studied. These are Families I, II, III, and VIII, the 
members of which are almost exclusively base-forming 
elements. The compounds of these elements with hy- 
drogen and oxygen are bases, or, in other words, have the 

427 



428 COLLEGE CHEMISTRY. 

power to neutralize acids. Their oxides are for the most 
part basic. An exception to this is found in the case of 
boron, already treated of, which forms a Aveak acid — boric 
acid. Its oxide is only slightly basic. The most strongly 
marked examples of base-forming elements are those which 
occur in Family I, Group A; then follow in order those 
of Group A, Family II, and Group A, Family III. The 
resemblance between the members of Group B, Family I, 
and those of Group A of the same family is less striking 
than the resemblance between the two groups of any other 
family. Between the members of Group B, Family II, 
and those of Group A of the same family there is a general 
resemblance, while there are also differences. A similar 
remark applies to the relations between Groups A and B, 
Family III. The members of Family VIII occupy a some- 
what exceptional position, as has already been pointed out. 
Each group of which this family consists is made up of 
three very similar elements with atomic weights which 
differ but little from one another. 

Metallic Properties. — It has long been customary to 
divide the elements into two classes — the metals and the 
non-metals. This classification was originally based upon 
differences in the physical properties of the elements, the 
name metal being applied to those elements which have 
what is known as a metallic lustre, are opaque, and are 
good conductors of heat and electricity. All those ele- 
ments which do not have these properties are called non- 
metals. Gradually the name metal came to signify an 
element which has the power to take the place of the 
hydrogen of acids and form salts, and the name non-metal 
to signify an element which has not this power. This 
classification, as will be seen, is practically the same as 
that which divides the elements into acid-forming and 
base-forming. The latter are the metals, the former are 
the non-metals. The imperfection of this classification 
has already been commented upon, the imperfection 
arising from the fact that some elements belong to both 
classes. 



STUDY OF THE BASE-FORMING ELEMENTS. 429 

Order in which the Base-forming Elements will be 
Taken Up. — In studying the base-forming elements, it 
appears best to begin with those which have the most 
strongly marked character. These are the members of 
Family I, Group A. It further appears best to adhere as 
closely as possible to the arrangement of the periodic 
system. Accordingly, the following order will be observed 
in the presentation of the elements yet to be studied: 

1. Elements of Family I, Group A, or the Potassium 
Group, consisting of lithium, sodium, potassium, ru- 
bidium, and caesium. 

2. Elements of Family II, Group A, or the Calcium 
Group, consisting of glucinum, magnesium, calcium, 
strontium, barium, and erbium. 

3. Elements of Family III, Group A, or the Aluminium 
Group, consisting of aluminium, scandium, yttrium, lan- 
thanum, and ytterbium. 

4. Elements of Family I, Group B, or the Copper 
Group, consisting of copper, silver, and gold. 

5. Elements of Family II, Group B, or the Zinc Group, 
consisting of zinc, cadmium, and mercury. 

6. Elements of Family III, Group B, or the Gallium 
Group, consisting of gallium, indium, and thallium. 

7. Elements of Family IV, Group B, or the Tin Group, 
consisting of germanium, tin, and lead. 

8. Elements of Family V, Group A, or the Vanadium 
Group, consisting of vanadium, columbium, didymium, 
and tantalum. 

9. Elements of Family VI, Group A, or the Chromium 
Group, consisting of chromium, molybdenum, tungsten, 
and uranium. 

10. Elements of Family VII, Group A, or the Man- 
ganese Group, of which manganese is the only representa- 
tive. 

11. Elements of Family VIII, of which there are three 
groups : 

(A) The Iron Group, consisting of iron, nickel, and 
cobalt; 



43° COLLEGE CHEMISTRY. 

(B) The Palladium Group, consisting of ruthenium, 
rhodium, and palladium; and 

(0) The Platinum Group, consisting of osmium, irid- 
ium, and platinum. 

Occurrence of the Metals. — One of the first questions 
that suggests itself in connection with each element is, In 
what forms of combination does it occur in nature ? The 
chemical compounds which occur ready-formed in nature 
are called minerals; and the minerals, and mixtures of 
minerals, from which the metals are extracted for practical 
purposes are called ores. The most common ores are oxides 
and sulphides. Examples of these are the ores of iron, 
tin, copper, lead, and zinc. The carbonates also occur in 
large quantity in nature, and are used for the purpose of 
preparing some of the metals. The carbonate of zinc, for 
example, is a valuable ore of zinc. 

Extraction of the Metals from their Ores. — The detailed 
study of the methods used in the extraction of the metals 
from their ores is the object of metallurgy. Besides the 
methods used on the large scale, there are others which 
are only used in the laboratory. The most common 
method of extracting metals from their ores is that used 
in the case of iron, which consists in heating the oxides 
with charcoal. If the ores used are not oxides, they must 
first be converted into oxides before this method is appli- 
cable. This can generally be accomplished by heatiug the 
ores in contact with the air. Under these circumstances 
the natural carbonates, sulphides, and hydroxides, are 
converted into oxides. These changes are illustrated by 
the following equations: 

FeCO s = FeO + C0 2 ; 
2FeO + = Fe 2 3 ; 
2FeS 2 + 110 = Fe 2 3 + 4S0 2 ; 
2Fe(OH) 3 = Fe 2 O s + 3H 2 0. 

By heating metallic oxides with carbon at the high tem- 
peratures attainable by means of powerful electric currents 



THE PROPERTIES OE THE MET4LS. 43 ' 

ill the electric furnace some oxides are easily reduced that 
cannot be reduced by other means. 

A second method consists in reducing the oxide by heat- 
ing it in a current of hydrogen. This has been illustrated 
in the action of hydrogen upon copper oxide, when the 
following reaction takes place: 

CuO + H 2 = H 2 + Ou. 

The method is efficient for many oxides, but is expensive 
and is not used on the large scale. 

Another method of extraction consists in treating the 
chloride of a metal with sodium. This is illustrated in 
the preparation of magnesium, which is made by heating- 
together magnesium chloride and sodium: 

MgCl 2 + 2Na = 2NaCl + Mg. 

Such a method is employed only in case it is impossible or 
extremely difficult to reduce the oxide. 

Besides the above methods, there are others which will 
be described under the individual metals. 

The Properties of the Metals. — As we shall find, the 
metals differ very markedly from one another. Some are 
light, floating on water, as lithium, sodium, etc. ; some are 
extremely heavy, as lead, platinum, etc. Some combine 
with oxygen with great energy; others form very unstable 
compounds with oxygen. Some form strong bases; others 
form weak bases. In general, those elements which are 
lightest, or which have the lowest sj^ecific gravity, are the 
most active chemically, while those which have the highest 
specific gravity are the least active. Among the former 
are lithium, sodium, and potassium; among the latter are 
lead, gold, and platinum. 

Compounds of the Metals.- — The principal compounds 
of the metals may be conveniently classified as: 

a. Compounds with fluorine, chlorine, bromine, and 
iodine; or the fluorides, chlorides, bromides, and iodides. 



43 2 COLLEGE CHEMISTRY. 

b. Compounds with oxygen, and with oxygen and 
hydrogen; or the oxides and hydroxides. 

c. Compounds with sulphur, and with sulphur and 
hydrogen; or the sulphides and hydrosulphides. 

d. Compounds with nitrogen; or the nitrides. 

e. Compounds with carbon and with silicon; or the 
carbides and. silicides. 

f. Compounds with the acids of nitrogen; or the nitrates 
and. nitrites. 

g. Compounds with the acids of chlorine, bromine, and 
iodine; or the chlorates, bromates, iodates, hypochlorites, 
etc. 

h. Compounds with the acids of sulphur, selenium, and 
tellurium; or the sulphates, sulphites, etc. 

i. Compounds with carbonic acid; or the carbonates. 

j. Compounds with the acids of phosphorus, arsenic, 
and antimony; or the phosphates, arsenates, etc. 

k. Compounds with silicic acid; or the silicates. 

I. Compounds with boric acid ; or the borates. 

It is more important to become acquainted with the 
general methods of preparation and the general properties 
of the more important compounds than to learn details 
concerning many individual members of each class. Only 
those compounds which illustrate general principles, or 
which, owing to some application, happen to be of special 
interest, need be dealt with in this book. 

The acids of which the salts are derivatives are already 
known t< us, and in dealing with the acids frequent refer- 
ence has been made to the methods of making the salts, 
and to some of their more important properties. It will 
be well, before taking up the metals systematically, briefly 
to treat of the general methods of preparation, and the 
general properties of the different classes of metallic com- 
pounds. It must be borne in mind, however, that the 
only way to become familiar with these substances and 
their relations is by working with them in the laboratory. 

Chlorides. — The chlorides, as well as the fluorides, 
bromides, and iodides, may be regarded as the salts of 



CHLORIDES. 433 

hydrochloric, hydrofluoric, hvdrobromic, and hydriodic 

acids, or simply as compounds of the metals with the 
members of the chlorine family. The most important of 
these compounds are the chlorides, and these well illus- 
trate the conduct of the others. 

The chlorides are made by treating a metal with chlo- 
rine, or with hydrochloric acid; by treating an oxide or a 
hydroxide with hydrochloric acid; by treating an oxide 
with chlorine and a reducing agent, like carbon: by treat- 
ing a salt of a volatile acid with hydrochloric acid; by 
treating a salt of an insoluble acid with hydrochloric acid; 
by adding hydrochloric acid or a soluble chloride to a solu- 
tion containing a metal with which chlorine forms an 
insoluble compound; and by adding to a solution of a 
chloride a salt, the acid of which forms with the metal of 
the chloride an insoluble salt, while the metal contained 
in it forms with chlorine a soluble chloride. 

Only two of the above methods are peculiar to chlorides. 
These are the treatment of the metals with chlorine, and 
the treatment of oxides with chlorine and a reducing 
agent. The others involve principles which are also in- 
volved in the preparation of all salts, and they may there- 
fore be treated of in a general way. 

The formation of chlorides by direct treatment of the 
metals with chlorine is the simplest method of all. It has 
been illustrated in studying chlorine. It was found that 
chlorine combines with other elements with great ease. 
Thus, iron, copper, and tin combine with it, as repre- 
sented in the following equations : 

Cu + CI = CuCl; 
2Fe + 3C1 2 = 2FeCl 3 ; 
JSn + 2C1 2 = SnCl 4 . 

The preparation of chlorides by treating oxides with 
chlorine and a reducing agent has been illustrated in the 
making of boron trichloride. It is used in making 
aluminium trichloride. For this purpose, chlorine is 



434 COLLEGE CHEMISTRY. 

passed over a heated mixture of aluminium oxide and 
charcoal, when reaction takes place according to the fol- 
lowing equation : 

A1 2 3 + 3C + 3C1 2 = 2A1C1 3 + 3CO. 

The other methods for preparing chlorides are, as has 
been said, general in character and are applicable to most 
salts. 

Formation of Salts in General. 

1. By treating a metal witli an acid. — This is the 
simplest method. It has been illustrated in the prepara- 
tion of zinc sulphate by the action of zinc on sulphuric 
acid • 

Zn + H 2 S0 4 = ZnS0 4 + H 2 . 

Other common examples are those represented in the fol- 
lowing equations : 

Fe + H 2 S0 4 = FeS0 4 + H 2 ; 
Zn + 2HC1 = ZnCl 2 + H 2 . 

2. By treating an oxide or a hydroxide with an acid. — 
This is of more general application than the preceding- 
method. As it has been studied in some detail in connec- 
tion with the subject of salts (see pp. 161-166), it need 
not be further treated of here. 

3. By treating the salt of a volatile acid with another 
acid. — This method has been repeatedly illustrated in the 
decomposition of carbonates and nitrites by acids in 
general. While carbonic acid and nitrous acid themselves 
are perhaps not formed in these reactions, and we cannot 
say that the carbonates and nitrites are salts of volatile 
acids, yet the decomposition products of these acids are 
volatile at ordinary temperatures. The decomposition of 
carbonates by acids has been pretty fully studied, though 
attention was not directed to the fact that this kind of 
action may be utilized for the purpose of making salts. 
As some carbonates occur in large quantity in nature or in 



FORMATION OF SALTS IN GENERAL 4 35 

the market, salts are frequently made by treating them 
with acids. Thus, magnesium sulphate is made by treat- 
ing magnesium carbonate with sulphuric acid: 

MgC0 3 + II 2 SO, = MgSO, + 11,0 + CO,; 

and calcium chloride is made by dissolving calcium car- 
bonate in hydrochloric acid : 

CaC0 3 + 2HC1 = CaCl a + H 2 + C0 2 ; etc. 

These reactions proceed to the end for the reason that one 
of the products is volatile and is carried out of the field 
of action (see p. 412). 

4. By treating a salt of an insoluble acid with another 
acid. — This case does not occur practically, as there are 
no common, insoluble acids. The principle involved is 
illustrated to some extent by the decomposition of a soluble 
silicate. Sodium silicate, for example, is soluble. When 
its solution in water is treated with an acid the silicic acid 
is partly precipitated, as we have seen : 

Na 2 Si0 3 + 2HC1 + H 2 = Si(OH) 4 + 2NaCl. 

The silicic acid formed is, however, not perfectly in- 
soluble in water, so that the reaction is not complete. In 
any case the reaction is not one that is used for the prep- 
aration of salts. 

5. By the action of tivo salts upon each other. — This 
method can be best described by means of an example. 
Suppose it is desired to prepare copper chloride by the 
action of two salts upon each other. Copper chloride is 
soluble. If copper sulphate and barium chloride are 
brought together in solution, the products are insoluble 
barium sulphate and soluble copper chloride: 

CuS0 4 + BaCl 2 = BaS0 4 + CuCl 2 . 

By simply filtering off the barium sulphate, a solution 
of copper chloride is obtained. 



43 6 COLLEGE CHEMISTRY. 

6. By premutation. — This method is illustrated in the 
formation of barium sulphate, referred to in the last para- 
graph. Obviously, it is applicable only to difficultly solu- 
ble or insoluble Lalts. Many carbonates and phosphates 
can be made in this way. 

General Properties of the Chlorides. — Most of the chlo- 
rides of the metals are soluble in water without decom- 
position, though many of them are decomposed when 
heated to a sufficiently high temperature with water. It 
will be remembered that the chlorides of the non-metallic 
or acid-forming elements are decomposed by water, yield- 
ing the corresponding oxides or hydroxides. The chlorides 
of some elements which are partly basic and partly acid 
are only partly decomposed. This is illustrated by the 
chloride of antimony, which with water forms an oxy- 
chloride : 

SbCl 3 + H 2 = SbOCl + 2HC1. 

The chlorides of the most strongly marked metals, like 
potassium, sodium, etc., are not decomposed by water. 
Calcium chloride dissolves with great ease, and, if the 
solution is evaporated, the chloride is again obtained. If, 
however, the attempt is made to drive off all the water by 
heat, some of the chloride is converted into the oxide as 
represented in the equation 

CaCl 2 + H 2 = CaO + 2HC1. 

Magnesium chloride is completely decomposed, if its solu- 
tion in water is evaporated to dryness, the action being 
the same in character as that which takes place in the case 
of calcium chloride. The chlorides of iron and aluminium 
and of many other metals act in the same way. Silver 
chloride and mercurous chloride, HgCl, are insoluble in 
water. Lead chloride is difficultly soluble in water. If, 
therefore, on adding hydrochloric acid or a soluble chlo- 
ride to a solution, a precipitate is formed, the conclusion 
is generally justified that one or more of the three metals 



THE SO-CALLED DOUBLE CHLORIDES. 437 

— silver, load, or mercury— is present. By taking into 
account the differences between these chlorides, it is not 
difficult to decide of which of them a precipitate consists. 
The chlorides are for the most part decomposed when 
treated with sulphuric acid, as has been shown in the 
action of sulphuric acid upon sodium chloride. Under 
these circumstances hydrochloric acid is given off, and the 
sulphate of the metal with which the chlorine was in com- 
bination is formed. In general, the reaction is represented 
by such equations as the following: 

2MC1 -f KSO i = M f S0 4 + 2HC1; 
M"Cl a + H 2 S0 4 = M"S0 4 + 2HC1; etc. 

Under ordinary circumstances, chlorides are not decom- 
posed by any acid except sulphuric acid. 

The So-called Double Chlorides and Similar Compounds 
of Fluorine, Bromine, and Iodine. — These compounds and 
their relations to the oxygen salts have been repeatedly 
referred to. Many chlorides combine with the chlorides 
of the stronger metals, like sodium and potassium, form- 
ing well-characterized compounds. Generally, these 
double chlorides are analogous to the oxygen salts in com- 
position, differing from them only by containing two 
atoms of chlorine in the place of each of the oxygen atoms. 
As examples of these salts of the chloro-acids those which 
are formed by the chlorides of platinum, antimony, 
chromium, and gold may be mentioned. Platinic chlo- 
ride, PtCl 4 , combines with other chlorides, forming salts 
of the general composition expressed by the formula 
PtCl 4 + 2MC1, or M 2 PtCl 6 . 

Different Chlorides of the Same Metal. — Just as sul- 
phur, selenium, phosphorus, and the other acid-forming 
elements combine with chlorine and the other members of 
the chlorine group in more than one proportion, so many 
of the metals combine with the members of the chlorine 
group in more than one proportion. Thus, mercury forms 
the two chlorides, HgCl 2 and HgCl, known respectively as 
mercuric and mercurous chlorides; iron forms the two 



43& COLLEGE CHEMISTRY. 

chlorides FeCl 3 and FeCl 2 , known as ferric and ferrouc 
chlorides; and tin forms stannic chloride, SnCl 4 , and 
stannous chloride, SnCl 2 . The conversion of a mgher 
chloride into a lower one is called an act of reduction. 
The change can generally be effected by means of nascent 
hydrogen : 

SnCl 4 + 2H = SnCl 2 + 2HC1. 
FeCl 3 + H = FeCl 2 + HC1. 

The conversion of a lower chloride into a higher one is 
generally spoken of as an act of oxidation, for the reason 
that it is most commonly effected by the action of oxygen. 
Thus the most convenient way to transform ferrous chlo- 
ride into ferric chloride is to treat it in solution in hydro- 
chloric acid with an oxidizing agent, when a double action 
takes place, as represented in the following equation : 

2FeCl 2 + 2HC1 + = 2FeCl 3 + H 2 0. 

The same change can be effected by the direct action of 
chlorine : 

FeCl a + 01 = FeCl 3 . 

In this case it would obviously be incorrect to speak of the 
process as one of oxidation. 

Another method of reduction, besides that referred to 
above, involving the action of nascent hydrogen, is that 
illustrated in the equation 

2HgCl 2 + SnCl 2 = 2HgCl + SnG\ 

In this case mercuric chloride is changed to mercurous 
chloride by the action of stannous chloride. The latter 
unites with chlorine so readily that it extracts it from 
some other chlorides, and is itself transformed into stannic 
chloride. While, therefore, we say that the stannous 
chloride reduces the mercuric chloride, it is equally true 
to say that the mercuric chloride chlorinates the stannous 
chloride. 



OXIDES. 439 

Oxides. — The oxides occur extensively in nature, and 
are among the most common ores of some of the impor- 
tant metals. The oxides of iron, tin, and manganese, for 
example, occur in nature. They can be made by oxidiz- 
ing the metals, by heating nitrates, carbonates, and 
hydroxides, and by heating some sulphides in contact with 
the air. 

When magnesium is burned it is converted into mag- 
nesium oxide: 

Mg + = MgO. 

When lead nitrate is heated it gives off oxygen and an 
oxide of nitrogen, and lead oxide is left behind: 

Pb(N0 3 ) 2 = PbO + 2NO, -f 0. 

When calcium carbonate is heated it yields calcium oxide 
and carbon dioxide : 

CaC0 3 = CaO + C0 2 

When aluminium hydroxide, A1(0H) 3 , is heated it loses 
water, and aluminium oxide is left behind : 

2A1(0H) S = A1 2 3 + 3H 2 0. 

The sulphide of iron, when heated in contact with the air, 
or "roasted," is converted into ferric oxide and sulphur 
dioxide. 

Most of the oxides of the metals are insoluble in water. 
Those of Group A, Family I, are soluble, but are con- 
verted by water into the corresponding hydroxides. 

The oxides are acted upon generally by acids forming 
the corresponding salts. If the salt with a certain acid is 
insoluble, the salt is not formed by the action of that acid 
on the oxide unless the acid or its anhydride is fusible and 
not volatile, when by fusing them together the salt is 
formed. 

Different Oxides of the Same Metal. — Just as there are 
different chlorides of the same metal, so there are different 



44° COLLEGE CHEMISTRY. 

oxides, and indeed there is greater variety among these 
than among the chlorides. Iron forms three oxides, ferric 
oxide, Fe 2 3 , ferroso-ferric oxide, Fe 3 4 , and ferrous 
oxide, FeO; mercury forms the two oxides HgO and Hg 2 ; 
etc. The lower oxides are converted into the higher by 
oxidation, and the higher into the lower by reduction. 
The higher oxides of several of the metals are acidic. 
This is markedly so in the case of chromium and man- 
ganese. 

Hydroxides. — The hydroxides are formed by treating 
oxides with water and by decomposing salts by adding 
soluble hydroxides to their solutions. In general, when- 
ever a salt is decomposed by a strong base, the base of the 
salt separates in the form of the hydroxide. The forma- 
tion of a hydroxide by the action of water on an oxide is 
well illustrated by the action of water on lime or calcium 
oxide, a process familiarly known as slaking : 

CaO + H 2 = Ca(OH) 2 . 

Most of the hydroxides of the metals are insoluble in 
water. If a soluble hydroxide is added to a solution con- 
taining a metal whose hydroxide is insoluble, the latter is 
precipitated. Thus, if a solution of sodium hydroxide is 
added to a solution of a magnesium salt, magnesium 
hydroxide is precipitated: 

MgSO, + SNaOH - Na 2 S0 4 + Mg(OH) 2 . 

So, also, when a solution of a ferric salt is treated with 
sodium hydroxide, a precipitate of ferric hydroxide is 
formed : 

FeCl 3 + 3NaOH = 3NaCl + Fe(OH) 3 . 

Only the hydroxides of the members of the potassium 
family, and some of the members of the calcium family, 
are soluble in water. The hydroxides of sodium and potas- 
sium are called alkalies. The solution of ammonia in 
water acts like a soluble hydroxide, and probably contains 



DECOMPOSITION OF SALTS BY BASES 441 

ammonium hydroxide, NH 4 (OH), formed by the action of 
water on ammonia : 

NH 3 + H 2 = NH 4 (OH). 

Now, when any one of the soluble hydroxides is added to 
a salt containing any metal that does not belong to the 
potassium or calcium family, an insoluble compound is 
formed. 

Decomposition of Salts by Bases. — The decomposition 
of salts by bases is analogous to the decomposition by 
acids. When a soluble base acts upon a salt, there are 
four possible kinds of action : 

1. The base from which the salt is derived may be vola- 
tile, or may break up, yielding a volatile product. 

In this case, decomposition takes place and the volatile 
base is given off. This is not a common case except among 
the compounds of carbon. The one illustration which 
we have had is the decomposition of ammonium salts by 
calcium hydroxide and sodium hydroxide, when the vola- 
tile compound ammonia, NH 3 , is given off. 

2. The hydroxide, or base from which the salt is 
derived, may be insoluble or difficultly soluble in water, 
and not volatile. 

In this case, if both the salt and the base are in solution, 
decomposition takes place, and the insoluble or difficultly 
soluble hydroxide, or base, is precipitated. This has 
already been illustrated. 

3. The base from which the salt is derived may be solu- 
ble and not volatile. 

This is the case, for example, when sodium hydroxide 
is added to a solution of potassium nitrate. Here sodium 
nitrate, potassium .nitrate, sodium hydroxide, and potas- 
sium hydroxide may all be present in the solution, and 
investigation has shown that all are present and that the 
quantity of each depends upon the masses of the sub- 
stances brought together, and upon their affinities. 

4. The fourth case is that in which a soluble hydroxide 



44 2 COLLEGE CHEMISTRY. 

forms an insoluble salt with the acid of a soluble salt, 
leaving a soluble hydroxide in solution. 

This is illustrated by the action of calcium hydroxide 
on a solution of sodium carbonate, when insoluble calcium 
carbonate is thrown down, and sodium hydroxide remains 
in solution, as represented in the equation 



Na a C0 8 + Ca(OH) 2 = 2NaOH + CaCO 



Some basic hydroxides, which are precipitated by solu- 
ble hydroxides, have a weak acid character, and, after they 
are precipitated, they redissolve in an excess of the soluble 
hydroxide. This is true, for example, of the hydroxides 
of aluminium, chromium, and lead. The salt-like com- 
pounds thus formed are generally quite unstable. Th« 
precipitation and subsequent solution of the hydroxides 
of the three metals named take place thus : 

AICI3 + 3NaOH = Al(OH), + 3NaCl; 
Al(OH), + 3NaOH = Al(ONa) 3 + 2H 2 0; 
CrCl 3 + 3NaOH = Cr(OH) 3 + 3NaCl; 
Cr(OH) 3 + 3NaOH = Cr(ONa) 3 + 3H 2 0; 
Pb(N0 2 ) 2 + 2NaOH = Pb(OH) 2 + 2NaN0 8 ; 
Pb(OH), + 2NaOH = Pb(ONa) 2 + 2H 2 0. 

In some cases where a soluble hydroxide is added to a 
salt, an oxide is precipitated instead of the hydroxide. 
This is analogous to the formation of an anhydride of an 
acid instead of the acid itself, as when carbonates, sul- 
phites, and nitrites are decomposed. 

When a silver salt is treated with a soluble hydroxide, 
silver oxide is at once precipitated. The same is true of 
mercury salts : 

2AgN0 3 + 2KOH = Ag 2 + H 2 -f- 2KN0 3 ; 
HgCl 2 -f 2NaOH = HgO + H 2 + 2NaCl. 

It is probable that the first product is the hydroxide, and 
that this breaks down into the oxide and water: 



SULPHIDES. 443 

2AgN0 3 + 2K0H = 2AgOH + 2KN0 3 ; 

2AgOH = Ag,0 + 11,0; 

H g Cl 2 + 2Na0H = Hg(0H) t + 2NaCl; 

Hg(0H) 2 = HgO + H 2 0. 

Some hydroxides are converted into the oxides by 
simply boiling the liquids in which they are suspended. 
Thus, when a salt of copper is treated with a soluble 
hydroxide, copper hydroxide is first precipitated; bat if 
the solution in which it is suspended is boiled, it is soon 
changed to the oxide : 

CuS0 4 + 2NaOH = Na 2 S0 4 + Cu(OH),; 
Cu(OH) 2 = 2 CuO + H a 0. 

The hydroxides corresponding to some of the higher 
oxides of the metals, as those of chromium and manganese, 
are acids. 

The hydroxides of most of the metals are decomposed 
by heat into water and the corresponding oxides. Those 
of the alkali metals, as potassium and sodium, are, how- 
ever, not decomposed by heat. 

Sulphides, — Many sulphides are found in nature as, for 
example, iron pyrites, FeS 2 ; lead sulphide, or galenite, 
PbS; copper pyrites, FeCuS 2 ; etc. They are made in the 
laboratory by heating metals with sulphur; by treating 
solutions of salts with hydrogen sulphide; by treating 
solutions of salts with soluble sulphides; and by reducing 
sulphates. Attention has been called to the fact that the 
sulphides are analogous in composition to the oxides, and 
that they are to be regarded as salts of hydrogen sulphide 
formed by substituting metals for the hydrogen of the acid. 

The formation of sulphides by the direct combination of 
sulphur with the metals is shown in the formation of lead 
sulphide and copper sulphide: 

Pb + S = PbS; 
2Cu + S = Cu 2 S. 



444 COLLEGE CHEMISTRY. 

The formation of sulphides by the action of hydrogen 
sulphide upon solutions of salts was discussed at some 
length under Hydrogen Sulphide (which see). The ex- 
tensive use made of this reaction in chemical analysis was 
also referred to. 

The action of soluble sulphides or solutions of salts is 
in general the same as that of hydrogen sulphide, but in 
some cases, in which the former will not act, the latter 
will. Thus, hydrogen sulphide will not precipitate iron 
sulphide from a solution of an iron salt, because iron sul- 
phide is easily acted upon by dilute acids. When, for 
example, hydrogen sulphide is passed into a solution of 
ferrous chloride, it naturally tends to form the sulphide 
FeS: 

FeCl, + H 2 S = FeS + 2HC1. 

But ferrous sulphide, FeS, is acted upon by dilute hydro- 
chloric acid, and is converted by it into ferrous chloride 
and hydrogen sulphide : 

FeS + 2HC1= FeCl 2 + H 2 S. 

It is therefore obvious that the first reaction cannot take 
place. 

If, however, a soluble sulphide, as sodium or ammonium 
sulphide, is added to a solution of an iron salt, iron sul- 
phide is precipitated, as in this case no free acid is formed. 
Thus, when ferrous chloride and ammonium sulphide are 
brought together the reaction takes place as represented 
in the equation 

FeCl 2 + (NHJ 2 S = FeS + 2NH 4 C1. 

In ammonium chloride the ferrous sulphide is nou 
soluble. 

The formation of a sulphide by reduction of a sulphate 
is illustrated by the formation of barium sulphide by heat- 
ing a mixture of barium sulphate and charcoal : 

. BaS0 4 + 4C = BaS + 4C0. 



HYDROSULPHIDFS. 445 

The sulphides of the alkali metals are soluble in water. 
Those of the other metals are insoluble. It should be 
remarked, however, that aluminium and chromium do not 
form sulphides, or, at least, if they do, the compounds are 
decomposed by water into hydroxides and hydrogen sul- 
phide. Barium sulphide is decomposed by water, and 
probably magnesium sulphide also. 

The sulphides are stable when heated without access of 
air; but if heated in the air they are converted into oxides 
of the metals and sulphur dioxide, or, in some cases, they 
take up oxygen and are converted into sulphates. The 
conversion of sulphides into oxides and sulphur dioxide 
by heating them in contact with the air has been repeatedly 
referred to. The process is carried on on the large scale 
in the j>reparation of iron ores for reduction, and is called 
roasting. The conversion of a sulphide into a sulphate 
by heating is a simple process of oxidation. Copper sul- 
phide is converted into the sulphate when heated for some 
time: 

CuS + 40 = CuS0 4 . 

This is the reverse of the reaction mentioned by which a 
sulphate is converted into a sulphide by reduction. 

Some sulphides, as those of sodium, potassium, and 
ammonium, take up sulphur in much the same way that 
they take up oxygen, and form the polysulphides. The 
two reactions appear to be analogous : 

K 5 S + 40 = K 2 S0 4 ; 

K 5 S + 4S =K 2 SS 4 , orK 2 S 5 . 

Hydrosulphides. — The hydrosulphides bear the same 
relation to the sulphides that the hydroxides bear to the 
oxides. They are not, however, as numerous nor as easily 
obtained as the hydroxides. 

The only hydrosulphides known are derived from the 
members of the potassium and calcium groups, and these 
are soluble. They are formed by saturating solutions of 



446 COLLEGE CHEMISTRY. 

the corresponding hydroxides with hydrogen sulphide. 
Potassium hydrosulphide is formed thus: 

KOH + H 2 S = KSH + H 2 0. 

Ammonium hydrosulphide is formed thus: 

NH 4 OH + H 2 S = NH 4 SH + H 2 0. 

It also appears probable that whenever a sulphide dissolves 
in water it is converted into a hydrosulphide and a hy- 
droxide. Thus it semes to be true that potassium sulphide 
is converted into the hydrosulphide and hydroxide : 

K 2 S + H 2 = KSH + KOH. 

Sulpho-salts. — The relation of the sulpho-salts to the 
sulphides has already been explained. It is like that of 
the ordinary oxygen salts to the oxides, and that of the 
chloro-salts, or double chlorides, to the chlorides. They 
are formed by dissolving the sulphides of certain metals, 
particularly tin, arsenic, and antimony, in the sulphides 
of the members of the potassium group : 

As 2 S 3 -f- 3K 2 S = 2K 3 AsS 3 ; 
As 2 S 5 + 3K 2 S = 2K 3 AsS 4 ; 
SnS 2 +K 2 S = K 2 SnS 3 , etc. 

These sulpho-salts are decomposed by the ordinary acids, 
the insoluble sulphides being precipitated thus : 

2K 3 AsS 3 + 6HC1 = As 2 S 3 + 6KC1 -f 3H 2 S- 
2K 3 AsS 4 + 6HC1 = As 2 S 5 + 6KC1 + 3H 2 S. 

Nitrates. — The nitrates are formed by dissolving the 
metals in nitric acid, and by treating oxides, hydroxides, 
carbonates, aud some other easily decomposed salts with 
nitric acid. The action of nitric acid upon metals was 
discussed under the head of Nitric Acid (which see). It 
was pointed out that the acid gives up a part of its oxygen 



CHLORATES. 447 

to the metal and forms different oxides, according to the 
conditions. Thus, when the acid acts upon copper the 
main product of the reduction is nitric oxide, but by 
changing the concentration of the acid a considerable 
quantity of nitrous oxide is formed. When zinc is dis- 
solved in nitric acid a part of the acid is reduced to 
ammonia. 

The nitrates are soluble in water, and all are decom- 
posed by heat. Some of them when heated lose only a 
third of their oxygen and are reduced to nitrites. This 
is true of potassium nitrate, the decomposition of which 
is represented by the equation 

KN0 3 = KN0 2 + 0. 

Most of the nitrates, however, are decomposed further, 
forming oxides. This has been shown in the case of lead 
nitrate, which when heated is converted into bad oxide, 
while nitrogen peroxide and oxygen are given off: 

Pb(N0 3 ) 2 = PbO + 2N0, + 0. 

If the oxide of the metal is decomposed by heat, as in the 
case of mercury, of course the product will be the metal. 

Chlorates. — These salts, except potassium chlorate, are 
not commonly met with. Potassium chlorate is manufac- 
tured in large quantity, and the other chlorates are 
generally made from it. The chlorates are soluble in 
water, and are decomposed by heat more easily than the 
nitrates are. They are first converted into perchlorates, 
and these are further decomposed by higher heat into 
chlorides and oxygen. 

The hypochlorites are formed by treating some of the 
metallic hydroxides in dilute solution with chlorine. This 
has been illustrated in the formation of " bleaching 
powder/' which contains calcium hypochlorite or a com- 
pound closely related to it. The hypochlorites, like the 
chlorates, are decomposed by heat. 



448 COLLEGE CHEMISTRY. 

Sulphates.— The general relations of the sulphates to 
sulphuric acid were treated of under Sulphuric Acid 
(which see). Some of these salts occur in nature in large 
quantity, as those of calcium and barium. The former is 
known as gypsum, the latter as heavy spar. Sulphates are 
made by treating metals, metallic hydroxides or oxides, 
carbonates, etc., with sulphuric acid; and by treating a 
solution containing a metal whose sulphate is insoluble, 
with sulphuric acid or a soluble sulphate. 

Zinc and iron give hydrogen and a sulphate when treated 
with sulphuric acid: 



Zn + H 2 S0 4 = ZnS0 4 '+ H 3 ; 
Fe + H 2 S0 4 = FeS0 4 + H 2 . 



This kind of, action takes place whenever a metal is dis- 
solved in sulphuric acid at the ordinary temperature. If, 
however, the temperature is raised the displaced hydrogen 
acts upon some of the sulphuric acid, or the metal extracts 
some of the oxygen of the acid, reducing it partly to sul- 
phurous acid, when sulphur dioxide is given off. This 
happens in the case of copper, as has been pointed out. 
It may be represented either by these equations: 

Cu + H 2 S0 4 = CuS0 4 + H 2 ; 
H 2 S0 4 + 2H = S0 2 + 2H 2 0; 

or by these : 

Cu +H 2 S0 4 = CuO -fS0 2 + H 2 0; 
CuO + H 2 S0 4 = CuS0 4 + H 2 0. 

The action of sulphuric acid on metallic hydroxides has 
been fully described. 

Most sulphates are soluble in water. The sulphates of 
barium, strontium, and lead are insoluble in water, and 
the sulphate of calcium is difficultly soluble. Therefore, 
if sulphuric acid or a soluble sulphate is added to a solu- 
tion containing either of the metals, barium, strontium, 
or lead, a precipitate is formed. A precipitate is also 



CARBONATES. 449 

formed when a concentrated solution o^ a calcium salt is 
treated in the same May. 

When heated with charcoal in the reducing flame of the 
blowpipe, sulphates are reduced to sulphides: 

K 2 S0 4 + 40 = K 2 S + 4C0, or 
K a S0 4 + 2C = K 2 S + 2C0 2 . 

Sulphites are made from sodium or potassium sulphite, 
which are made by treating sodium or potassium hydroxide 
in solution with sulphur dioxide: 

2NaOH + S0 2 = Na 2 S0 2 + H 2 0. 

All sulphites are decomposed by the common acids, sul- 
phur dioxide being given off: 

Na 2 SO, + H 2 S0 4 = Na,S0 4 + H 2 -f S0 2 . 

The sulphites are changed to sulphates by oxidation. 
Thus, sodium sulphite is changed to the sulphate when its 
solution is allowed to stand in contact with the air : 

Na 3 S0 8 + = Na 2 S0 4 . 

The sulphites, like the sulphates, are reduced to sul- 
phides. 

Carbonates. — Many carbonates are found in nature, 
some of them in great abundance and widely distributed. 
The principal one is calcium carbonate. They are made 
by passing carbon dioxide into solutions of hydroxides, 
and by adding soluble carbonates to solutions of salts con- 
taining metals whose carbonates are insoluble. 

The formation of carbonates by the action of carbon 
dioxide on a solution of a hydroxide is illustrated in the 
case of potassium hydroxide : 

2K0H + C0 2 = K 2 C0 3 + H 2 0. 



45° COLLEGE CHEMISTRY. 

The formation of calcium carbonate takes place in the 
same way, but the carbonate formed is insoluble : 

Ca(OH), + C0 2 = CaC0 3 + H 2 0. 

If in either case the action is continued, the norma 
carbonate first formed is converted into acid carbonate: 



K 2 C0 3 + C0 2 + H 2 = 2KHCO 

Oa j 
HJ 



& 2 C0 3 -f C0 2 + H 2 5= 2KJd 
CaC0 3 + C0 2 + H 2 = Oa | {CQ ^ 



All normal carbonates except those of the members of 
the potassium family are insoluble, and are decomposed 
by heat into carbon dioxide and the oxide of the metal. 
The decomposition of calcium carbonate into lime and 
carbon dioxide is an example : 

CaC0 3 = CaO + C0 2 . 

When a soluble carbonate is added to a solution of a 
calcium, barium, or strontium salt, the corresponding in- 
soluble carbonates are precipitated. When a magnesium 
salt is treated with a soluble carbonate, however, a basic 
carbonate is precipitated. 

Phosphates. — Calcium phosphate is very abundant in 
nature, and a few other phosphates are also found. The 
methods used for making phosphates are the same as those 
used in making salts in general. 

The normal phosphates of all the metals except the 
members of the potassium family are insoluble in water. 
The normal phosphates, as a rule, are not changed by 
heat. The secondary phosphates, such as secondary 
sodium phosphate, HNa 2 P0 4 , lose water when heated, and 
yield pyrophosphates: 

2HNa 2 P0 4 = Na 4 P 2 7 -f H 2 0. 

Sodium 
. pyrophosphate. 



SILICATES. 45 I 

Those phosphates in which only one-third of the 
hydrogen is replaced by metal — as, for example, primary 
sodium phosphate, H 2 NaP0 4 — lose water when heated, 
and yield metaphosphates : 

H 2 NaP0 4 = NaP0 3 + H 2 0. 

Sodium 
metaphosphate. 

» 

Neither the pyrophosphates nor the metaphosphates are 
changed by heat. 

Silicates. — The silicates, as has been stated, are very 
widely distributed in nature. Those which are most 
abundant are the feldspars and their decomposition- 
products. The principal feldspar is a complex silicate of 
aluminium and potassium, of the formula KAlSi 3 8 , 
derived from the polysilicic acid H 4 Si 3 8 , which is formed 
from three molecules of normal silicic acid by the loss of 
four molecules of water: 

3Si(OH) 4 = H 4 Si 3 8 + 4H 2 0. 

Silicates can be made by heating together, at a high 
temperature, silicon dioxide, in the form of sand, and 
basic oxides or carbonates: 

CaO + Si0 2 = CaSi0 3 ; 
Na 3 CO s + Si0 2 = Na 2 Si0 3 + C0 2 . 

Only the silicates of the members of the potassium group 
are soluble in water. When these are treated in solution 
with dilute acids, they are decomposed, as has been 
explained under Silicic Acid (which see). 

Some silicates, which are insoluble in water, are decom- 
posed by the ordinary acids, such as sulphuric and hydro- 
chloric acids, the silicic acid separating as a difficultly 
soluble substance, which, if dried on the water-bath, 
becomes insoluble. 

Many silicates, which are not acted upon by strong 
acids, are decomposed by fusion with sodium or potassium 



452 COLLEGE CHEMISTRY. 

carbonate, when the silicate of potassium or sodium and 
the oxide of the metal contained in the silicate are formed. 
Silicates which are not decomposed in either of the ways 
mentioned yield to hydrofluoric acid. The action con- 
sists in the formation of the gas, silicon tetrafluoride, 
SiF 4 , and the fluorides of the metals present. Thus, the 
reaction in the case of feldspar takes place in accordance 
with the equation, 

KAlSi 3 8 -j- 16HF = KF + A1F 3 + 3SiF 4 + 8H 2 0. 

The silicon fluoride is given off as a gas, and the fluorides 
formed are soluble in water. Hence, hydrofluoric acid is 
said to dissolve the silicates. 



EXPERIMENTS. 

Chlorides, Bromides, and Iodides. 

Experiment 201.— Dissolve a small crystal of silver nitrate in 
pure water. Add to a small quantity of this solution in a test- 
tube a few drops of dilute hydrochloric acid. The white sub- 
stance thus precipitated is silver chloride, AgCl. To another 
small portion of the solution add a few drops of a dilute solution 
of common salt, or sodium chloride, NaCl. The white substance 
produced in this case is also silver chloride. Add ammonia to 
each tube. If sufficient is added the precipitates will dissolve. 
On adding enough hydrochloric acid to these solutions to com- 
bine with all the ammonia the silver chloride is again thrown 
down. On standing exposed to the light both precipitates change 
color, becoming finally dark violet. The reactions involved in the 
above experiments are these : In the first place, when hydro- 
chloric acid is added to silver nitrate this reaction takes place : 

Ag + N0 3 4- H + CI = AgCl + H + N0 3 . 
When sodium chloride is added this reaction takes place : 

+ - 4- + 

Ag 4- N0 3 + Na + CI = AgCl + Na + N0 3 . 



EXPERIMENTS : HYDROXIDES. 453 

In the first reaction the ions of nitric acid are formed ; in tho 
second, the sodium and silver exchange places. In addition to 
the insoluble silver chloride, there are formed at the same time 
the ions of the soluble salt, sodium nitrate. On adding ammonia 
the silver chloride forms with it a compound which is soluble in 
water ; and on adding an acid, the ammonia combines with it, 
leaving the silver chloride uncombined and therefore insoluble. 

Extensive use is made of insoluble compounds for the purpose 
of detecting substances in analysis. The only insoluble chlorides 
are those of silver, lead, and mercury.* If, therefore, on adding 
hydrochloric acid or a soluble chloride to a solution, a precipitate 
is formed, the conclusion is justified that one or more of the three 
metals — silver, lead, or mercury — is present. By taking account 
of the differences in the properties of these chlorides it is not dif- 
ficult to decide of which of them a precipitate consists. 



Hydroxides. 

Experiment 202. — To some pieces of freshly-burnt lime add 
enough cold water to cover it. The action which takes place is 
represented by the equation 

CaO + H a O = Ca(OH) 2 . 

The process is known as slaking. 

Experiment 203. — To a small quantity of a dilute solution of 
magnesium sulphate add a dilute solution of caustic soda. The 
white precipitate is magnesium hydroxide. [Would you expect 
this precipitate to be soluble in sulphuric acid? in hydrochloric 
acid ? in nitric acid ?] The answers follow from these considera- 
tions : When acids act upon hydroxides, salts are formed ; mag- 
nesium sulphate is soluble, as is seen by the fact that we started 
with a solution of this salt ; the only insoluble chlorides are those 
of silver, lead, and mercury ; all nitrates are soluble. 

When a solution of an iron salt is treated with sodium hydrox- 
ide a precipitate of iron hydroxide is formed : 

FeCl, + 3NaOH = Fe0 3 H 3 + 3NaCl. 

Experiment 204. — To a dilute solution of that chloride of iron 
which is known as ferric chloride add caustic soda. The reddish 

* There are two chlorides of mercury. Only one of them, mer- 
curous chloride, is insoluble. 



454 COLLEGE CHEMISTRY. 

precipitate which is formed is ferric hydroxide. [From the gen- 
eral statements made above, would you expect this precipitate to 
be soluble in hydrochloric acid ? in nitric acid ? Try each. Is it 
soluble in sulphuric acid ?] 

Experiment 205. — Add to a solution of an aluminium salt 
sodium hydroxide. After a precipitate is formed continue to add 
the sodium hydroxide. Perform similar experiments with a 
chromium and with a lead salt. Boil each of the solutions ob- 
tained. Treat a solution of copper sulphate with sodium hydrox- 
ide in the cold. Heat. . 



Sulphates. 

Experiment 206.— Make a dilute solution of barium chloride, 
of lead nitrate, of strontium nitrate. To a small quantity of each 
in a test-tube add a little sulphuric acid. [What remains in solu- 
tion?] Make a somewhat concentrated solution of calcium chlo- 
ride. To this add sulphuric acid. [What is in solution ?] Add 
more water, and see whether the precipitate will dissolve. The 
formulas of the salts used in the experiments are barium chloride, 
BaCl a ; lead nitrate, Pb(N0 3 ) 2 ; strontium nitrate, Sr(N0 3 ) 2 . 
[Write the equations expressing the reactions.] If to the solutions 
of the salts any soluble sulphate is added instead of sulphuric acid, 
the same insoluble sulphates will be formed. The sulphates of iron, 
copper, sodium, and potassium are among the soluble sulphates. 
Make dilute solutions of small quantities of each of these, and add 
them successively to the solutions of barium chloride, lead nitrate, 
and strontium nitrate. The formula of iron sulphate is FeS0 4 ; of 
copper sulphate, CuSO* ; of sodium sulphate, Na 2 S0 4 ; and of po- 
tassium sulphate, K 2 S0 4 . Write the equations representing the 
reactions which take place in the above experiments. It need 
hardly be explained that the action consists in an exchange of 
places on the part of the metals. Thus, when the soluble salt iron 
sulphate, FeS0 4 , is brought together with the soluble salt barium 
chloride, BaCl 2 , the insoluble salt barium sulphate, BaS0 4 , and 
the soluble salt iron chloride, FeCl 2 , are formed : 

FeS0 4 + BaCl 2 == FeCl 2 + BaS0 4 . 



Reduction of Sulphates to Sulphides. 

Experiment 207. — Mix and moisten a little sodium sulphate 
and finely-powdered charcoal. Heat the mixture for some time 
in the reducing flame. After cooling scrape off the salt, dissolve 



EXPERIMENTS: CARBONATES. 455 

it in a few cubic centimetres of water, and filter through a small 
filter. If the change to the sulphide has taken place, sodium sul- 
phide, Na 2 S, is in solution. A solution of a sulphide when added 
to a solution containing copper gives a black precipitate of copper 
sulphide. Try this ; also try the action on the solution of the 
salt of copper of some of the sulphate from which the sulphide 
was made. 

Carbonates. 

Experiment 208.— The formation of carbonates by the addi- 
tion of soluble carbonates to solutions of salts of metals whose 
carbonates are insoluble, is illustrated by the following experi- 
ments : Make solutions of copper sulphate, iron sulphate, lead 
nitrate, silver nitrate, calcium chloride, barium chloride. Add 
to each a little of a solution of a soluble carbonate, as sodium 
carbonate, potassium carbonate, ammonium carbonate. Note 
the result in each case. Filter off all the precipitates and prove 
that they are carbonates. This may be done by treating them 
with dilute acids, which decompose them, causing an evolution 
of carbon dioxide, which can be detected by passing a little of it 
into lime-water. In some of the cases mentioned the insoluble 
salts formed are basic carbonates, as, for example, those of cop- 
per and magnesium. The salts of silver, calcium, and barium 
are the normal carbonates Ag 2 C0 3 , BaC0 3 , and CaC0 3 . 



CHAPTER XXV. 

ELEMENTS OF FAMILY I, GROUP A: 
THE ALKALI METALS :— LITHIUM.— SODIUM.— POTAS- 
SIUM.— RUBIDIUM— CAESIUM.— AMMONIUM. 

General. — The elements of this group which are most 
abundant in nature are sodium and potassium. While 
lithium occurs in considerable quantity, the two remaining 
elements, rubidium and caesium, have been found in only 
very small quantities. They are all strongly basic, their 
hydroxides being the strongest bases known. They form 
well-characterized salts with all acids, and as a rule their 
salts are very stable. In all their compounds they act as 
univalent elements, except in those which they form with 
hydrogen, and in their peroxides .; in the latter they appear 
to be bivalent. Leaving these compounds out of consid- 
eration the general formulas of some of the other principal 
compounds are as follows: 

MCI, M 2 0, M 2 S, M(OH), M(SH), MNO„ M 2 S0 4 , etc. 

The valence of the members of the group towards other 
elements is, in general, constant. 

The relations between the atomic weights are interest- 
ing. That of sodium, 23.05, is very nearly half the sum 
of those of lithium, 7.03, and potassium, 39.15. We have 

7 ' 03 + 39 - 15 = 23.09. 

So, also, that of rubidium, 85.4, is approximately half the 
sum of those of potassium, 39.15, and caesium, 133. 

39.15 -f 133 

~ = 86.08. 

456 



POTASSIUM. 457 

The specific gravity of these elements increases with the 
atomic weight; and their melting-points become lower as 
the atomic weights become higher. 

At. Wt. Sp. Grav. M. P. 

Lithium 7.03 0.594 180°. 

Sodium 23.05 0.972 95.6° 

Potassium 39.15 0.865 02.5° 

Rubidium 85.4 1.52 38.5° 

Caesium 133 ? ? 

The regularity is complete in the case of the melting- 
points, but as regards the specific gravities sodium is an 
exception to the rule. As sodium and potassium and 
their compounds are much more commonly met with than 
the other members of the group, these will form the chief 
subject of this chapter. 

Potassium, K (At. Wt. 39.15). 

Occurrence. — Potassium is a constituent of many min- 
erals, particularly of feldspar, the common variety of 
which, as has already been explained, is a complex silicate 
of aluminium and potassium. It is found also in com- 
bination with chlorine as carnallite and sylvite; with sul- 
phuric acid and aluminium, as alum ; with nitric acid, as 
saltpetre or potassium nitrate ; and in other forms. The 
natural decomposition of minerals containing potassium 
gives rise to the presence of this metal in various forms of 
combination everywhere in the soil. It is taken up by the 
plants; and, when vegetable material is burned, the potas- 
sium remains behind, chiefly as potassium carbonate. 
When wood-ash is treated with water the potassium car- 
bonate is dissolved, and it can be obtained in an impure 
state by evaporating the solution. The substance thus 
obtained is called potash. In the juice of the grape there 
is contained a salt of potassium, mono-potassium tartrate, 
which is deposited in large quantity from wine. This is 
commonly called "crude tartar." 



45 g COLLEGE CHEMISTRY. 

Preparation. — Potassium was first prepared by Davy in 
the year 1807, by the action of a powerful electric current 
on potassium hydroxide. It was then for many years pre- 
pared by heating to a high temperature a mixture of 
potassium carbonate and carbon, obtained by heating in a 
closed vessel ordinary mono-potassium tartrate deposited 
from wine. 

K 2 C0 3 + 2C = 2K + 3CO. 

The metal is now prepared by electrolysis of potassium 
hydroxide or potassium chloride. 

Properties. — Potassium is a light substance, which floats 
on water. Its freshly cut surface has a bright metallic 
lustre, almost white; it acts energetically upon water, 
causing the evolution of hydrogen, which, together with 
some of the potassium, burns, while potassium hydroxide 
is formed at the same time. This reaction has been 
studied in connection with hydrogen. In consequence of 
its action upon water, potassium cannot be kept in the 
air. It is kept under some oil, as petroleum, upon which 
it does not act. In an atmosphere upon which it does not 
act, as, for example, hydrogen, it can be distilled. Its 
vapor is green. Its specific gravity is 0.865; its melting- 
point 62.5°. It combines with chlorine and bromine with 
great energy, and has the power to extract chlorine from 
its compounds. It can, therefore, be used for the purpose 
of isolating some elements, as, for example, magnesium 
and aluminium, whose oxygen compounds cannot be 
reduced by the ordinary methods. As, however, sodium 
is generally used for this purpose instead of potassium, on 
account of its lower price, the action will be referred to 
more at length under Sodium. Although the metal is 
converted into vapor, no reliable determination of the 
specific gravity of the vapor has been made, for the reason 
that the vessels which have been used for the purpose have 
always been acted upon, and the results thus vitiated. 

Potassium Hydride, K 2 H. — This compound is formed 
by heating potassium in an atmosphere of hydrogen at 



POTASSIUM SALTS. 459 

about 300°. It is a silver-white mass with a metallic 
Lustre. It takes fire in the air. When heated, it begins 
to dissociate at 200°. 

(Fluoride, KF 

Potassium -j B ro mide KBr — ^ these salts the only 
[iodide, KI 

one that occurs in nature in quantity is the chloride. 
This is found in the great salt deposits at Stassfurt, 
Germany, and in some other localities in the form of the 
mineral sylvite, which is more or less impure potassium 
chloride. It is also found in the form of a compound 
containing magnesium, potassium, and chlorine, of the 
formula MgCl 2 .KCl + GH 2 0, or KMgCl 3 + 6H ? 0, known 
as carnallite. 

The other salts of the group are made by the general 
methods for making salts, that is, by neutralizing the acids 
with the hydroxide or carbonate of potassium. It is, 
however, easier to make the iodide by other methods, and 
as there is a large demand for this salt for use in medicine 
and in the art of photography, several methods have been 
devised for its preparation. Of these, only one need be 
described. This consists in treating iron filings under 
water with iodine. Both the iron and iodine dissolve, 
forming ferrous iodide, Fel 2 . If to the solution of this 
compound half as much iodine is added as has already been 
used in its preparation, ferroso-ferric iodide, Fe 3 I 8 , is 
formed and remains in solution. By adding a solution of 
potassium carbonate to this, reaction takes place as repre- 
sented in the equation: 

Fe 3 I 8 + 4K 2 C0 3 + 4II 2 = 8KI + Fe s (OH) 8 + 4C0 2 . 

The hydroxide of iron is insoluble, and can be removed 
by filtration. 

The fact that the specific gravity of hydrofluoric acid at 
a low temperature corresponds to the formula H 2 F 2 makes 
it not improbable that potassium fluoride has the formula 
K 3 F 2 . This appears still more probable from the fact that 



460 COLLEGE CHEMISTRY. 

there is an acid potassium fluoride of the formula KHF 2 , 
or KF -J- HF. Similar acid salts have not been obtained 
from the other acids of the group. 

Properties. — All these salts are soluble, and crystallize 
well in cubes. The fluoride is the most easily soluble in 
water. If deposited from a water solution at the ordinary 
temperature the crystals contain two molecules of water of 
crystallization, and are deliquescent. The iodide is solu- 
ble in 0.7 part of water at the ordinary temperature, and 
is also soluble in alcohol (40 parts). The bromide requires 
about 1^ parts of water for solution at the ordinary tem- 
perature, and is but slightly soluble in alcohol. The 
chloride is soluble in 3 parts of water at the ordinary tem- 
perature, and is insoluble in alcohol. All are decom- 
posed by sulphuric acid. 

Applications. — Potassium chloride is extensively used 
for the purpose of making other potassium salts, as, for 
example, the nitrate and carbonate; the bromide is used 
in medicine; the iodide, as stated above, is used in medi- 
cine and in photography. 

Potassium Hydroxide, KOH. — This well-known sub- 
stance, commonly called caustic potash, is prepared by 
treating potassium carbonate in solution with calcium 
hydroxide in a silver or iron vessel. The reaction is based 
upon the fact that calcium carbonate is insoluble, and that 
potassium carbonate and calcium hydroxide are soluble : 

K 2 C0 3 + Ca(OH) 2 = 2KOH + CaC0 3 . 

Potassium hydroxide is also prepared on the large scale 
by the action of the electric current on a solution of 
potassium chloride: 

2KC1 + 2H 2 = 2KOH + H 2 + Cl 2 . 

It is a white brittle substance. In contact with the air 
it deliquesces, and absorbs carbon dioxide, being com- 
pletely transformed into potassium carbonate. It is the 
strongest of the bases. It decomposes the salts of all other 



POTASSIUM OXIDE. 461 

bases, even of those which, like sodium and lithium 
hydroxides, are soluble in water. 

Animal substances like the skin are disintegrated by the 
hydroxide. It has a caustic action. 

Instead of potassium hydroxide, the corresponding 
sodium compound is used wherever this is possible, as the 
latter is cheaper. The chief application of the potassium 
compound outside of the laboratory is for making soft- 
soap. For this purpose fats are boiled with a solution of 
potassium hydroxide or carbonate. 

Potassium Oxide, K 2 0.— This compound can be made 
by burning potassium in the air, and heating the residue 
to a high temperature. It is also formed by melting 
potassium hydroxide and metallic potassium together: 

2K + 2KOH - 2K 2 + H 2 . 

"With water it forms the hydroxide, with a marked evolu- 
tion of heat: 

K,0 + H 2 = 2KOH. 

Potassium also forms other oxides of which the peroxide 
of the formula K 2 + is the best studied. This peroxide is 
the final product of the combustion of potassium in the air 
or in oxygen. At a high temperature it breaks down into 
potassium oxide, K 2 0, and oxygen. It also gives up its 
oxygen very readily to substances which are capable of 
oxidation, acting so energetically upon some as to cause 
evolution of light. 

Potassium Hydrosulphide, KSH, is analogous to potas- 
sium hydroxide. Just as the latter is made by the action 
of potassium on water, so the former can be made by the 
action of potassium on hydrogen sulphide : 

K 2 + 2H 2 S = 2KSH + H 2 . 

It is, however, obtained most readily by the action of 
hydrogen sulphide on a solution of potassium hydroxide: 

KOH + H 2 S = KSH + H 2 0. 



462 COLLEGE CHEMISTRY. 

Potassium Nitrate, KN0 3 . — This salt is commonly 
called saltpetre. Its occurrence in nature has already been 
referred to under Nitric Acid (which see). When refuse 
animal matter is left to undergo decomposition in the 
presence of bases, nitrates are always the end-products. 
They are consequently found very widely distributed in 
the soil. In the East Indies the potassium nitrate formed 
in the neighborhood of dwellings and stables is collected, 
and sent into the market. The process of nitrification is 
carried on artificially on the large scale in the so-called 
'•'saltpetre plantations/' In these, refuse animal matter 
is mixed with earthy material, wood ashes, etc. , and piled 
up. These piles are moistened with the liquid products 
from stables. After the action has continued for two or 
three years the outer crust is taken off, and extracted with 
water. The solution thus obtained contains, besides potas- 
sium nitrate, calcium and magnesium nitrates. It is 
treated with a water-extract of wood ashes or with potas- 
sium carbonate, by which the calcium and magnesium are 
thrown down as carbonates. Much of the saltpetre in the 
market is made from Chili saltpetre, or sodium nitrate, by 
treating it with potassium chloride, advantage being taken 
of the fact that sodium chloride is less soluble in hot water 
than potassium nitrate. Molecular weights of sodium 
nitrate and potassium chloride are dissolved in water and 
the solution heated. When the solution is sufficiently 
concentrated, sodium chloride is deposited: 

NaN0 3 + KC1 = KN0 3 + NaCl. 

This being removed, potassium nitrate crystallizes from 
the mother-liquors. 

Potassium nitrate crystallizes in long rhombic prisms, 
of a salty taste. Under some circumstances it crystallizes 
in rhombohedrons. When dissolved in water it causes a 
lowering of the temperature. At ordinary temperatures 
1-00 parts of water dissolve from 20 to 30 parts of the salt; 
at 100°, 100 parts dissolve 247 parts. 



GUNPOIVDBR. 4^3 

Applications. — Potassium nitrate is used as an oxidizing 
agent in the laboratory, and in the manufacture of fire- 
works. Its chief use, however, is in the manufacture of 
gunpowder. 

Gunpowder. — The value of gunpowder is due to the fact 
that it explodes readily, the explosion being a chemical 
change accompanied by a sudden evolution of gases. It 
is a mixture of saltpetre, charcoal, and sulphur. When 
heated, the saltpetre gives oif oxygen and nitrogen; the 
oxygen combines with the charcoal, forming carbon dioxide 
and carbon monoxide; and the sulphur combines with the 
potassium, forming potassium sulphide. When a mixture 
of saltpetre and charcoal is burned, the reaction which 
takes place is this : 

■4KNO3 + 50 = 3C0 2 + 2N 2 + 2K 2 C0 3 . 

By adding the necessary quantity of sulphur the carbon 
dioxide, which would otherwise remain in combination 
with the potassium as potassium carbonate, is given off, 
and potassium sulphide formed: 

2KN0 3 + 3C + S = 3C0 2 + N a + K 2 S. 

For this reaction the constituents should be mixed in the 
proportions: 

Saltpetre 74.83 

Charcoal 13.31 

Sulphur 11.86 

100.00 

This is approximately the composition of all powder. 
When gunpowder explodes, the gases formed occupy about 
280 times the volume occupied by the powder itself. 

Potassium Nitrite, KX0 2 , is formed simply by heating 
the nitrate to a sufficiently high temperature. The reduc- 
tion is, however, much facilitated by adding to the nitrate 
some easily oxidized metal, as lead or iron. When the 
gases formed by the action of arsenic trioxide on nitric 



464 COLLEGE CHEMISTRY. 

acid are passed into potassium hydroxide, both the nitrite 
and nitrate are formed, and they can be separated by 
crystallization. 

Potassium Chlorate, KC10 3 . — The character of the 
reaction by which potassium chlorate is formed when 
chlorine acts upon a solution of potassium hydroxide has 
already been discussed (see p. 147). In the manufacture 
of the chlorate it is found advantageous to make calcium 
chlorate, and then to treat this with potassium chloride, 
when, at the proper concentration, potassium chlorate 
crystallizes out, on account of the fact that it is less solu- 
ble than the salts which are brought together. The 
process in brief consists in passing chlorine into a solution 
of calcium hydroxide in which an excess of hydroxide is 
held in suspension. The first action consists in the 
formation of calcium hypochlorite. When the solution 
of this salt is boiled it is decomposed, yielding the chlorate 
and chloride: 

3Ca(OCl) 2 = Ca(0 3 Cl) 2 + 2CaCl 2 . 

On now treating the solution with potassium chloride the 
following reaction takes place: 

Ca(0 3 Cl) a + 2KC1 = 2KOIO3 + CaCl 2 . 

Potassium chlorate crystallizes in lustrous crystals of 
the monoclinic system. Its taste is somewhat like that of 
saltpetre. It melts at a comparatively low temperature 
(334°), and at 352° begins to decompose, with evolution 
of oxygen. At ordinary temperatures 100 parts of water 
dissolve 6 parts of the salt, and at the boiling temperature 
60 parts. In consequence of the ease with which it gives 
up its oxygen, the chlorate is an excellent oxidizing agent, 
and it is constantly used in this capacity in the laboratory. 
Its oxidizing action is well illustrated by grinding a very 
little of it in a mortar with a little sulphur,* when an 

* Great care should be taken with all experiments with potassium 
chlorate. See description of experiments, 



POTASSIUM CYANIDE. 465 

explosion takes place. With phosphorus the action is 
exceedingly violent. 

The chief uses of potassium chlorate are in the prepara- 
tion of oxygen, and in the manufacture of matches and fire- 
works. The tips of Swedish safety matches are made of 
potassium chlorate and antimony sulphide. The surface 
upon which they are rubbed to ignite them contains red 
phosphorus. The chlorate is frequently used in medicine, 
particularly as a gargle in cases of sore throat. 

Potassium Perchlorate, KC10 4 , is formed in the first 
stage of the decomposition of the chlorate by heat, as was 
explained under Oxygen (which see). It is prepared best 
by heating the chlorate in an open vessel until, after 
having been liquid, it begins to get solid again. 

Potassium Cyanide, KCN. — Under Cyanogen it was 
stated that when nitrogen is passed over a highly-heated 
mixture of carbon and potassium carbonate, potassium 
cyanide is formed; and that carbon containing nitrogen 
compounds, as animal charcoal, when ignited with potas- 
sium carbonate, reduces the carbonate, forming potassium, 
in presence of which the carbon and nitrogen combine, 
forming the cyanide. The simplest way to make the 
cyanide is by heating potassium ferrocyanide, K 4 Fe (CN) 6 , 
with potassium : 

K 4 Fe(CN) 6 + 2K = 6KCN + Fe. 

Potassium cyanide is extremely easily soluble in water, 
and is deliquescent in moist air. When boiled with water 
it is decomposed, forming potassium formate and am- 
monia : 

KCN + 2H 2 = KC0 2 H + NH 3 . 

It combines readily with oxygen when in the molten con- 
dition, as shown in its action upon lead oxide (see p. 
380). In consequence of this power to combine with 
oxygen to form the cyanate, it is a valuable reducing 
agent, and is not infrequently used in the laboratory in 
this capacity. 



466 COLLEGE CHEMISTRY. 

Potassium cyanide combines with the cyanides of the 
metallic elements, forming the so-called double cyanides 
of which potassium ferrocyanide, K 4 Fe(CN) 6 , is a good 
example. 

Potassium Sulphocyanate, KCNS. — Just as potassium 
cyanide takes up oxygen to form the cyanate, it also takes 
up sulphur to form the sulphocyanate : 

KCN + S = KONS. 

It is easily prepared by adding sulphur to molten potas- 
sium cyanide, or by heating a mixture of dehydrated 
potassium ferrocyanide, potassium carbonate, and sulphur. 
It crystallizes particularly well out of its solution in 
alcohol. It is deliquescent, and when dissolved in water 
it causes a very considerable lowering of the temperature. 
Thus, when 500 grams of the salt are mixed with 400 
grams of water at the ordinary temperature, the tempera- 
ture sinks to about — 20°. Unlike the cyanate, it is not 
decomposed by water. 

Potassium Sulphate, K 2 S0 4 . — This salt occurs in com- 
bination with others in nature, particularly in the mineral 
hainite, which contains the constituents of potassium sul- 
phate, magnesium sulphate, and magnesium chloride, as 
indicated in the formula K 2 S0 4 .MgS0 4 .MgCl 2 + 6H 2 0. 
This occurs in Stassfurt and in Kalusz. Potassium sul- 
phate is used in medicine, and in the preparation of 
ordinary alum and of potassium carbonate. 

Primary, or Acid, Potassium Sulphate, KHS0 4 . — This 
salt is obtained as a secondary product in the preparation 
of nitric acid by the action of sulphuric acid upon salt- 
petre. It occurs in nature in the Grotto del Solfo, near 
Naples ; and is made by treating the neutral salt with con- 
centrated sulphuric acid. When heated above its melting- 
point it gives off water, and is transformed into the 
disulphate, thus: 

2KHS0 4 = K 2 S,0 7 + H 2 0. 

When the disulphate is heated in contact with basic oxides 



POTASSIUM SULPHITES— CARBONATES. 467 

it breaks down, forming sulphates. The decomposition is 
that represented in the equation 

K 2 S 2 0, = K.SO, + SO,. 

The nascent sulphur trioxide thus set free acts with great 
energy upon the oxides which are present. Hence acid 
potassium sulphate is a valuable agent for the purpose of 
decomposing some mineral substances which do not readily 
yield to the ordinary reagents. Its action consists in 
breaking down into the disulphate and water, the disul- 
phate then further breaking down into normal sulphate 
and sulphur trioxide. 

Sulphites. — When sulphur dioxide is passed into a solu- 
tion of potassium carbonate until carbon dioxide ceases to 
escape, potassium sulphite, K 2 S0 3 , is formed. If the gas 
is passed to saturation the product is the primary or acid 
sulphite, KHSO3. 

Carbonates. — The normal salt, K 2 C0 3 , is the chief con- 
stituent of wood-ashes. When these are extracted with 
water the carbonate passes into solution and the salt thus 
obtained can be purified in a number of ways. The impure 
salt is known as potash. Formerly all the potassium car- 
bonate made was obtained from wood-ashes, but at present 
not more than half of the supply comes from this source. 
The other sources are the residues from the manufacture 
of beet-sugar, potassium sulphate and chloride, and wool- 
fat. The preparation of the carbonate from the sulphate 
and chloride is accomplished by the same method as that 
used in the preparation of sodium carbonate from the 
chloride. The methods used for this purpose will be 
treated of under the head of Sodium Carbonate (which 
see). The salt crystallizes from very concentrated solu- 
tions in water. It is deliquescent, and dissolves in water 
with an evolution of heat, and the solution has a strong 
alkaline reaction. 

The explanation of the alkaline reaction of potassium 
carbonate in terms of the theory of electrolytic dissociation 



468 COLLEGE CHEMISTRY. 

is very interesting. When the salt is dissolved in water it 
is dissociated into ions thus : 

K 2 C0 3 = 2K + C0 3 . 
With water, which is itself partly dissociated, the ions 

C0 3 combine to some extent to form the ions HC0 3 , while 
at the same time hydroxyl ions are left in the solution as 
shown in the equation : 

C0 3 + H 2 = HOO3 + OH. 

But hydroxyl ions are believed to be the cause of basic 

reactions, so that the solution shows an alkaline reaction. 

It will be observed that the cause of the phenomenon is 

what we call the weakness of the carbonic acid. A weak 

acid is only slightly dissociated in water solution. If the 

anions of the acid are formed in the presence of water 

they unite with hydrogen ions to form undissociated acid 

molecules, and this necessitates the formation of hydroxyl 

ions. These latter, however, do not combine with the 

+ 
cations K which are in the solution, as the product of this 

union would be potassium carbonate which is dissociated 
by the water. 

Acid Potassium Carbonate, HKC0 3 , is formed by pass- 
ing carbon dioxide over the normal salt, or into the con- 
centrated aqueous solution of the latter. It is much less 
easily soluble in water than the normal salt. The dry salt 
gives off carbon dioxide and water easily when heated, and 
is converted into the normal salt : 

2KHC0 3 = K 2 C0 3 + C0 2 + H 2 0. 

The same decomposition takes place when the water solu- 
tion is heated, and even on evaporation at the ordinary 
temperature. 

Phosphates. — Three phosphates of potassium are known : 
(1) Tertiary , or normal potassium phosphate, K 3 P0 4 ; (2) 



RUBIDIUM— C/FSIUM. 4^9 

Secondary, or di-potassium phosphate, K 2 HP0 4 ; and (3) 
primary ', or mono-potassium phosphate, KH 2 P0 4 . There 

is nothing particularly characteristic about these salts, 
except the decompositions which the primary and secon- 
dary salts undergo when heated. These decompositions 
have already been referred to (see p. 450 and p. 451). 

Potassium Silicate, K 2 SiO s . — A compound of the definite 
composition represented by the formula here given has not 
been prepared. A solution of potassium silicate in water 
is prepared by dissolving sand or amorphous silicon dioxide 
in potassium carbonate or hydroxide. It is prepared on 
the large scale by melting together quartz powder and 
purified potash. It is known as water glass, for the reason 
that its solution dries in the air, forming a glass-like look- 
ing mass. To distinguish it from the water glass made 
with a sodium carbonate or hydroxide it is called potash 
water glass. 



Rubidium, Eb (At. Wt. 85.4). 
Caesium, Cs (At. Wt. 133). 

Both these elements are widely distributed, but only in 
small quantities. They generally occur in company with 
potassium, which they resemble closely. They were dis- 
covered by means of the spectroscope by Bunsen and 
Kirchhoff. The characteristic spectrum of rubidium con- 
sists of two dark red lines, and this is the origin of the 
name rubidium (from rubidus, dark red). Caesium was 
found in the Durkheim mineral water, and was recognized 
by two characteristic blue lines, and the name caesium was 
given to it on this account (from ccesius, sky-blue). 
Rubidium is found in different varieties of mica, known 
as lepidolite. The mineral pollitx, which is essentially a 
silicate of caesium and aluminium, contains caesium as one 
of the chief constituents. 

It is a remarkable fact that the elements rubidium and 
caesium which are so similar to potasssium accompany it 
so generally in nature. Similar facts were noted in the 



4)o COLLEGE CHEMISTRY. 

group consisting of chlorine, bromine, and iodine, and 
that of sulphur, selenium, and tellurium. It will he 
remembered that chlorine is frequently accompanied by 
bromine and iodine; and sulphur by selenium and tel- 
lurium; but that chlorine and sulphur are present in much 
larger quantities than the elements which accompany 
them. Further, the relations between the atomic weights 
of the members of each group are approximately the same. 

Rubidium is prepared by the same method as that used 
in the preparation of potassium. 

It is silver- white with a yellowish tint. It can be con- 
verted into vapor which has a blue color. It takes fire in 
the air at the ordinary temperature. Its action upon 
water is the same as that of potassium, and its salts are 
very similar to those of potassium. 

Caesium has been prepared by the electrolysis of molten 
caesium cyanide. 

The salts of caesium are much like those of rubidium 
and potassium. 



Sodium, Na (At. Wt. 23.05). 

Occurrence, — Sodium occurs very widely distributed and 
in large quantities in nature, principally as sodium chlo- 
ride. It is found in a number of silicates, and is a con- 
stituent of plants, especially of those which grow in the 
neighborhood of the seashore and in the sea. Just as the 
ashes of inland plants are rich in potassium carbonate, so 
the ashes of sea-plants and those which grow near the sea 
are rich in sodium carbonate. It is found everywhere in 
the soil, but generally in small quantities. Its presence 
in the soil is due to the decomposition of minerals con- 
taining it, such as soda feldspar, or albite. It occurs also 
as sodium nitrate or Chili saltpetre, and in large quantity 
in Greenland in the form of cryolite, which, as has been 
explained, is a so-called double "fluoride of aluminium and 
sodium, of the formula Na,AlF f , or AlF„.3NaF. 

3 6 ' o 



SODIUM. 471 

Preparation. — It is prepared by the action of an elec- 
tric current upon sodium hydroxide, or upon sodium 
chloride. 

Properties. — The properties of sodium are very similar 
to those of jotassium. It is light, floating on water; it 
has a bright metallic lustre; and at the ordinary tempera- 
ture it is soft like wax. It decomposes water, but not as 
readily as potassium does. Its specific gravity is 0.9735; 
its melting-point 95.6°. Its vapor is colorless when seen 
in thin layers, while thick layers appear purple. When 
melted and allowed to cool it takes the crystallized form. 
When exposed to the air it acts upon the moisture, and is 
converted into the hydroxide. 

Applications. — It is used for the purpose of isolating 
some elements whose oxides cannot easily be reduced, as, 
for example, aluminium, magnesium, and silicon, which 
are prepared by treating their chlorides with sodium. 
Silicon, however, as we have seen, is prepared better by 
treating potassium fluosilicate, K 2 SiF 6 , with sodium. 
The element is also used, in combination with mercury as 
sodium amalgam, a substance which affords a ready means 
of making nascent hydrogen. It also finds constant appli- 
cation in the laboratory for a variety of purposes. 

Sodium Hydride, Na 2 H, is formed in the same way as 
the corresponding compound of potassium, and is in every 
way similar to it. 

Sodium Chloride, NaCl. — This is the substance which 
is generally known simply as salt, or common salt. It 
occurs very widely distributed, and in immense quantities, 
in the earth. The most important deposits are those at 
Wieliczka in Galicia, at Stassfurt and Reichenhall in 
Germany, and at Cheshire in England. Besides these 
there are, however, many other deposits in the United 
States of America, in Africa, and in Asia. As it is easily 
soluble in water, many springs and streams, as well as 
lakes and the ocean, contain it. Sea-water contains 2.7 
per cent. In some places sodium chloride is taken out of 
mines in solid form. Frequently, however, water is 



472 COLLEGE CHEMISTRY. 

allowed to flow into cavities in the earth, and to remain^ 
for some time in contact with the salt. The solution thus 
formed -is afterward drawn or pumped out of the mine and 
evaporated by appropriate methods. It is generally allowed 
slowly to run down walls made of twigs, so that a large 
surface of the liquid is exposed to the air. The concen- 
trated solution thus obtained is then evaporated to crystal- 
lization by the aid of heat. 

In hot countries salt is obtained by the evaporation of 
sea-water, the heat of the sun being used for the purpose. 
Large shallow cavities are made in the earth, and into 
these the water flows at high-tide, or it is pumped up into 
them if they are too high. The process is continued for 
some months, and then the mother-liquor is drawn off, 
and the accumulated salt collected and subjected to proper 
methods of purification. 

The salt obtained by the above methods is not pure. It 
always contains sodium sulphate, together with magnesium 
and calcium chlorides. The chlorides of magnesium and 
calcium being deliquescent cause it to become moist in the 
air. Pure salt is not deliquescent. 

Sodium chloride crystallizes in colorless, transparent 
cubes. Some of that which occurs in nature has a blue 
color. When deposited from an evaporating solution it 
takes the form of small cubes arranged in groups of the 
shape of hollow pyramids, known as the hopper-shaped 
deposits. When deposited, the crystals enclose water, not 
as water of crystallization, and this is given off when the 
crystals are heated, the action being accompanied by a 
crackling sound. This is known as decrepitation. 

Sodium chloride melts at 815°, and is volatile at a red 
heat. In hot water it is but little more soluble than in 
cold. At 100° 100 parts of water dissolve 39 parts, and 
at ordinary temperatures 36 parts. 

Sodium chloride is the starting-point in the preparation 
of all sodium compounds, as well as of chlorine and 
hydrochloric acid. Salt is necessary to the life of man 
and many other animals. The role played by it in the 



SODIUM COMPOUNDS. 473 

animal economy is not understood, but it is found generally 
distributed throughout the body in small quantity. 

The fluoride, bromide, and iodide of sodium are like the 
corresponding potassium salts and need not be described. 

Sodium Hydroxide, NaOH. — This compound resembles 
potassium hydroxide in all respects. Being cheaper it is 
used more extensively. It is prepared in the same way, 
by treating sodium carbonate in solution with calcium 
hydroxide, when insoluble calcium carbonate and soluble 
sodium h) T droxide are formed: 

Na 2 C0 3 + Ca(OH) 2 = CaC0 3 + 2NaOH. 

It is now prepared, for the most part, by the electrolysis 
of sodium chloride. 

The substance is commonly called caustic soda. It is 
extensively used for the purpose of making soap from fats. 

Oxides. — Sodium forms two oxides, the monoxide, 
Na 2 0, and the peroxide, Na 2 2 . In this respect a differ- 
ence is noticed between sodium and potassium; the latter 
forming the compounds K 2 and K 2 4 . 

Sodium Peroxide, Na 2 2 , has acquired importance in 
the arts as a bleaching-agent. It is prepared by heating 
sodium in a current of dry air at a temperature of 300°. 
When heated to a high temperature it gives off oxygen. 
Water decomposes it, forming sodium hydroxide, and 
setting oxygen free. 

Sodium Sulphantimonate, Na 3 SbS 4 , also known as 
Schlippe's salt, is a particularly beautiful example of the 
salts of sulpho-acids. It is made, as its composition indi- 
cates, by dissolving antimony pentasulphide in a solution 
of sodium sulphide : 

Sb 2 S. + -3Na a S = 2Na 3 SbS 4 . 

Sodium Nitrate, NaN0 3 .— This compound occurs in 
large quantity in southern Peru on the border of Chili, 
and is known as Chili saltpetre. The natural salt con- 
tains, besides the nitrate, sodium chloride, sulphate, and 



474 COLLEGE CHEMISTRY. 

iodate. Sodium nitrate is similar to potassium nitrate, 
but it cannot be used in place of the more expensive potas- 
sium salt in the manufacture of the finer grades of gun- 
powder, as it becomes moist in the air, and does not 
decompose quickly enough. It is used extensively in the 
manufacture of nitric acid, and also for the purpose of 
preparing ordinary saltpetre. The iodine contained in 
Chili saltpetre is now extracted on the large scale, and this 
forms an important source of iodine. 

Sodium Sulphate, Na 2 S0 4 . — This salt was first made by 
Glauber, as it is now made, by the action of sulphuric acid 
on sodium chloride. It is commonly called Glauber's salt. 
It occurs in a number of natural waters, as in that of 
Friedrichshall and Carlsbad. It occurs, further, in solid 
form in small quantities in some localities. It is made in 
very large quantities in connection with the manufacture 
of soda, the first reaction in this process being that of 
sodium chloride upon sulphuric acid. It is also formed 
in the manufacture of nitric acid by the action of sulphuric 
acid on Chili saltpetre. 

It crystallizes in large, colorless, monoclinic prisms, 
which contain ten molecules of water. These crystals are 
formed, however, only in case the temperature of the 
solution is below 20° at the time they are deposited. The 
salt is most easily soluble in water at 33°; above this point 
the solubility decreases. If deposited from its solution at 
temperatures above 33° it is anhydrous. 

Sodium sulphate easily forms supersaturated solutions 
which crystallize rapidly if disturbed, if a small crystal of 
the salt is thrown into them, and if cooled down to — 8°. 
This phenomenon is frequently presented by salts, but it 
is shown in a particularly striking way by this one. 

When exposed to the air the salt loses its water of 
crystallization and crumbles to a white powder. This is 
the process already described as efflorescence (see p. 84). 

Sodium sulphate is used as a purgative in medicine, and 
in the laboratory for the production of cold artificially. 
A good freezing mixture is made by bringing it together 



SODIUM CARBONATE. 475 

with concentrated hydrochloric acid. Sodium chloride is 
formed, and the water of crystallization of the sulphate 
takes the liquid form. This change from the solid to the 
liquid form is accompanied by a marked absorption of 
heat. Ice can be made in this way without difficulty. 
The chief uses of the sulphate are in the manufacture of 
sodium carbonate and of glass, as will be explained farther 
on. 

Sodium Thiosulphate, ]STa 2 S 2 3 + 5li 2 0.— This is the 
salt which is commonly called hyposulphite of soda. It is 
made on the large scale by treating caustic soda with sul- 
phur, and conducting sulphur dioxide into the solution. 
It is also made by boiling a solution of sodium sulphite 
and adding sulphur: 

Na 2 S0 3 + S = Na 2 S 2 3 . 

Its chief application is in photography, in which art it is 
used for the purpose of dissolving the excess of silver salt 
on the plate which has been exposed to the light, and on 
which a picture has been developed. The action consists 
in the formation of salts in which both sodium and silver 
are contained. These are soluble in water. 

Sodium Carbonate, Na 2 C0 3 . — This salt, commonly called 
soda, is one of the most important of manufactured chem- 
ical compounds. The mere mention of the fact that it is 
essential to the manufacture of glass and soap will serve 
to give some conception of its importance. It is found in 
the ashes of sea-plants, just as potassium carbonate is 
found in the ashes of inland plants. Formerly, it was 
made entirely from plant-ashes, but we are no longer 
dependent upon this source for our supply of the salt, as 
two methods have been devised for preparing it from 
sodium chloride, with which nature provides us in such 
abundance. As these methods are of great importance, 
and are, further, very interesting applications of chemical 
principles, they will be described below. 

Properties. — Anhydrous sodium carbonate is a powder 
which is formed by heating the crystallized salt. It melts 



476 COLLEGE CHEMISTRY. 

to a clear liquid when heated to a sufficiently high tem- 
perature. It dissolves in water very readily with evolution 
of heat. The action is, however, not as marked as in the 
case of potassium carbonate. When the salt is deposited 
from a water solution it has the composition Na 2 C0 3 -f- 
10H 2 O. This salt, it will be observed, contains the same 
number of molecules of water of crystallization as sodium 
sulphate. Like this, too, it effloresces when exposed to 
the air. When heated it melts in its water of crystalliza- 
tion, and the salt Na 2 C0 3 + H 2 0, or (HO) 2 C(ONa) 2 , 
separates. This, however, loses water when heated higher, 
and is converted into the anhydrous salt. The conduct 
of the carbonate towards water at different temperatures 
is suggestive of that of the sulphate. Its maximum solu- 
bility is at temperatures between 33° and 70°. Above the 
latter point the solubility decreases. The crystals of 
sodium carbonate containing ten molecules of water of 
crystallization belong to the monoclinic system. 

Applications. — Sodium carbonate is used in immense 
quantities in the manufacture of glass, and in the prepara- 
tion of caustic .soda, which is used in the manufacture of 
soap. 

The Le Blanc Process for the Manufacture of Sodium 
Carbonate. — In the manufacture of soda the problem to 
be solved is the conversion of sodium chloride into sodium 
carbonate. The first method devised for this purpose is 
that of Le Blanc. During the French Kevolution the 
supply of potash was cut off from France. This led the 
government to offer a prize for a practical method for 
manufacturing soda from common salt. The method 
proposed by Le Blanc at that time, and which, until 
recently, has been used almost exclusively involves three 
reactions : 

(1) The sodium chloride is converted into sodium sul- 
phate by treating it with sulphuric acid : 

2Na01 + H 2 SO. = Na 2 S0 4 + 2HC1. 



SODIUM CARBONATE: LE BLANC PROCESS. 477 

(2) The sodium sulphate thus obtained is heated with 
charcoal, which reduces it to sodium sulphide : 

Na 2 S0 4 + 2C = Na a S + 2C0 2 . 

(3) The sodium sulphide is heated with calcium car- 
bonate, when sodium carbonate and calcium sulphide are 
formed : 

Na 2 S + CaC0 3 = Na 2 €0 3 + CaS. 

The conversion of the sulphate into the carbonate is, 
therefore, expressed by the equation 

Na 2 S0 4 + 2C + CaC0 3 = Na 3 C0 3 + CaS + 2C0 3 . 

Calcium sulphide is insoluble in water containing lime, 
so that by treating the resulting mass with water the 
sodium carbonate is separated from the sulphide. 

In practice the sodium sulphate is mixed with coal and 
calcium carbonate, and the mixture heated in appropriately 
constructed furnaces. The coal reduces the sulphate to 
sulphide, which then reacts upon the calcium carbonate 
as above represented. The product of the action is known 
as crude soda or black ash. It contains, as its chief con- 
stituents, sodium carbonate and calcium sulphide, together 
with some calcium oxide, and a number of other substances 
in small quantities. In order to purify this product, it is 
broken to pieces, and treated with water; and the solution 
thus obtained evaporated, when the salt of the composition 
Na 2 C0 3 -+- 2H 2 is deposited. This is dipped out, and 
dried by heat, when it loses all its water. The product is 
the calcined purified soda of commerce. This always con- 
tains some sulphate and chloride togther with a small 
quantity of sulphite. 

When dissolved in water and allowed to crystallize, the 
salt is deposited in large crystals which contain water in 
the proportion represented by the formula Na 2 C0 3 -f- 
10H 2 O. This is the so-called crystallized soda. 



47 8 COLLEGE CHEMISTRY. 

As, in the manufacture of soda, by the Le Blanc 
process, the sulphur remains in combination as calcium 
sulphide, a process, known as the Chance process, has 
been devised for its recovery. This consists in passing 
carbon dioxide into the waste, thus liberating hydrogen 
sulphide; passing this into another portion of the waste, 
thus converting the calcium sulphide into the hydrosul- 
phide; and then treating this with carbon dioxide, when 
a gas rich in hydrogen sulphide is given off: 

C0 2 + CaIi 2 S 2 + H 2 = CaC0 3 + 2H 2 S. 

By regulating the supply of air the gas is burned either to 
sulphur dioxide or to sulphur. 

Ammonia Process for the Manufacture of Soda. — 

Another process now in extensive use for the manufacture 
of soda is the so-called ammonia process, or the Solvay 
process. This depends upon the fact that mono-sodium 
carbonate, HNaC0 3 , is comparatively difficultly soluble in 
water. If, therefore, mono-ammonium carbonate, or acid 
ammonium carbonate, HNH 4 C0 3 , is added to a solution 
of common salt, acid sodium carbonate, HNaC0 3 , crystal- 
lizes out, and ammonium chloride remains in the solution: 

NaCl + HNH 4 C0 3 = HNaC0 3 + NH 4 C1. 

When the acid carbonate thus obtained is heated, it gives 
off carbon dioxide, and is converted into the normal salt 
thus : 

2HNaC0 3 - Na 2 C0 3 + C0 2 + H 2 0. 

The carbon dioxide given off is passed into ammonia, and 
thus again obtained in the form of acid ammonium car- 
bonate : 

NH 3 + H 2 + C0 2 = HNH 4 C0 3 . 

The ammonium chloride obtained in the first reaction 
is treated with lime or magnesia, MgO, and the ammonia 
set free. This ammonia is used again in the preparation 
of acid ammonium carbonate. The object of using mag 



SODIUM PHOSPHATES. 479 

nesia is to get magnesium chloride, which, when evaporated 
to dryness and heated, yields magnesia and hydrochloric 
acid : 

MgCl 2 -f H 2 = MgO + 2HC1. 

A large proportion of the soda supply of the world is now 
furnished by the Solvay process. 

Mono-Sodium Carbonate, Primary Sodium Carbonate, 
HNaCO s . — This salt is commonly called " bi-carbonate of 
soda." It is easily prepared by passing carbon dioxide 
over the ordinary carbonate dissolved in its water of crys- 
tallization : 

Na 2 C0 3 + C0 2 + H 2 = 2HNaC0 8 . 

"When heated it gives up carbon dioxide and water, and is 
converted into the normal salt. As was stated in connec- 
tion with the ammonia-soda process, primary sodium 
carbonate is much more difficultly soluble in water than 
the normal salt. At ordinary temperatures 100 parts of 
water dissolve about 10 parts of the salt. 

It is used in medicine, and extensively in the prepara- 
tion of ir soda water" and other effervescing drinks. 

Phosphates. — There are three phosphates of sodium 
just as there are three phosphates of potassium. The 
point of chief interest presented by them is that the 
secondary salt, HNa 2 P0 4 , is the one most easily obtained, 
and is the substance commonly known as sodium phos- 
phate. When a solution of this salt is treated with an 
excess of sodium hydroxide, and the solution evaporated, 
normal or tertiary sodium phosphate crystallizes out. This 
has the composition Na 3 P0 4 -f 12H 2 0. The solution of 
the latter salt has an alkaline reaction, and when exposed 
to the air it absorbs carbon dioxide, and is converted into 
the secondary salt: 

2Na s P0 4 + C0 2 + H 2 = 2HNa 2 P0 4 + Na 2 C0 3 . 

Secondary sodium phosphate, HNa 2 P0 4 -J- 12H,0, is 
easily made by adding sodium carbonate to a solution of 



4S0 COLLEGE CHEMISTRY. 

phosphoric acid until an alkaline reaction is reached. It 
is also prepared on the large scale from bone-ash. It 
forms monoclinic prisms which effloresce in the air. 

Sodium Borate. — Normal boric acid, as we have seen, 
has the composition B(OII) 3 , and there are a number of 
borates derived from this acid. The salt which boric acid 
most readily forms with sodium hydroxide or sodium car- 
bonate, however, is that derived from ietrdboric acid, 
H 2 B 4 7 , which is derived from normal boric acid by elim- 
ination of water. (See p. 342. ) This salt is borax, which 
in crystallized form has the composition represented by 
the formula Na 2 B 4 7 + 10H 2 O. 

Borax occurs in nature in several lakes in Asia and in 
Clear Lake, Nevada, in the United States. It is manu- 
factured by neutralizing, with sodium carbonate, the boric 
acid found in Tuscany. When heated, borax puffs up, 
and at red heat it melts, forming a transparent, colorless 
liquid. The dehydrated salt is known as anhydrous or 
calcined borax. In the molten condition, borax has the 
power to combine with metallic oxides, and, as many of 
the double borates thus formed are colored, the salt is used 
in blowpipe work for the purpose of detecting certain 
metals. As it dissolves metallic oxides, it is used in the 
process of soldering, as it is necessary to have bright, un- 
tarnished metallic surfaces in order that the solder shall 
adhere firmly. The action of molten borax upon metallic 
oxides is similar to that which takes place when sodium 
hydroxide acts upon a solution of borax. Borates of the 
metals are formed together with sodium borate, or double 
borates in which part of the hydrogen is replaced by 
sodium and part by other metals. 

Borax is extensively used in the manufacture of porcelain 
and in glass-painting. It is an antiseptic, preventing the 
decomposition of some organic substances. 

Sodium Silicate, Na 2 SiO s . — Sodium silicate is formed 
by dissolving silicon dioxide in sodium hydroxide, and 
can be obtained in crystallized form. It is prepared on 
the large scale by melting together quartz sand and sodium 



LITHIUM-AMMONIUM SALTS. 48 1 

carbonate in the proper proportions, and by melting 
together sodium sulphate, quartz sand, and charcoal 
powder. This substance is commonly known as water- 
glass. It is soluble in water, and, when its solution dries, 
it leaves a transparent coating on the surface on which it 
is placed. It is extensively used in the manufacture of 
artificial stone. 



Lithium, Li (At. Wt. 7.03). 

Lithium occurs in nature in relatively small quantity, 
chiefly in the minerals lepidolite, petalite, and spodumene, 
and in many mineral waters. It is also found in the ashes 
of a number of plants. It is prepared by the electrolysis 
of the chloride in the molten condition. The metal is 
silver-white, and is characterized by its low specific gravity. 
It acts vigorously upon water, but, if the water is at the 
ordinary temperature, the hydrogen given off does not 
take fire. In the air it conducts itself in much the same 
way that sodium does. 

The most characteristic salts of lithium are the phos- 
phate, carbonate, and chloride. 



Ammonium Salts. 

Attention has already been called to the marked simi- 
larity of the salts of potassium and sodium to those formed 
by the action of ammonia on the acids, and known as 
ammonium salts. The most important of these salts will 
be briefly treated of in this connection. A characteristic 
property of ammonium salts which distinguishes them 
from the salts of all the metals is their volatility. When 
sublimed, they all undergo decomposition, which is either 
partial or complete. 

The simplest kind of decomposition which they undergo 
is-dissociation into ammonia and the acid. This is illus- 
trated in the case of ammonium chloride, which, when 



482 COLLEGE CHEMISTRY. 

heated to a sufficiently high temperature, is dissociated 
into ammonia and hydrochloric acid. This is an example 
of true dissociation. The amount of decomposition is 
constant for any given temperature and pressure. 

An ammonium salt of a polybasic acid containing some 
metal gives off ammonia and leaves an acid salt, which 
generally undergoes further decomposition. Thus, sodium- 
ammonium sulphate, NaNH 4 S0 4 , first gives off ammonia 
and forms mono-sodium sulphate: 

s °.jra 1 = S0 <!H a + ^ 

The acid salt thus formed then undergoes further 
change and the pyrosulphate is formed : 

2S0 4 j ^ a = Na 2 S 2 7 + H 2 0. 

Another example of this kind of decomposition of am- 
monium salts is that afforded by sodium-ammonium phos- 
phate, HNaNH 4 P0 4 . When heated, this gives off ammonia 
and then water, the final product being sodium metaphos- 
phate : 

ONa ( (Ma 

PO { (MH 4 = PO \ OH + NH„; 
OH (OH 

ONa 
PO \ OH = P0 2 ONa + H 2 0. 
OH 

Some ammonium salts undergo deeper-seated decom- 
positions, and do not give ammonia as one of the products. 
This is true especially of such salts as readily give off 
oxygen. In such cases the ammonia is oxidized, so that 
the hydrogen forms water. This is illustrated in the 
decomposition of ammonium nitrate and nitrite : 

NH 4 N0 3 = N 2 + 2H 2 0; and 
NH^O, = N 2 + 2H 2 0. 



AMMONIUM CHLORIDE— AMMONIUM SULPHIDE. 4%3 

Further, all ammonium salts are decomposed with evolu- 
tion of ammonia when treated with basic hydroxides. 
This has been illustrated in the preparation of ammonia 
from ammonium chloride by treatment with calcium 
hydroxide : 

2NH 4 C1 + Ca(OH) 3 = CaCl 3 + 2NH 3 + 2H 3 0. 

The ammonium salts are made by neutralizing acids with 
ammonia. 

Ammonium Chloride, NH 4 C1. — This salt is commonly 
called sal ammoniac. At present its principal source is 
the so-called ammoniacal liquor of the gas-works. This 
liquid contains a considerable quantity of ammonium car- 
bonate, and, when it is treated with lime, ammonia is 
given off. This is passed into hydrochloric acid, and the 
solution of ammonium chloride thus formed evaporated to 
crystallization. The salt has a sharp, salty taste, and is 
easily soluble in water. When heated, it is converted into 
vapor without melting and with very slight decomposition ; 
and when the vapor comes in contact with a cold surface, 
it condenses in crystalline form. This process of vaporiz- 
ing and condensing a solid is called sublimation. Some 
of the ammonium chloride met with in the market has 
been sublimed. The salt is used in the preparation of 
ammonia, in medicine, and for other purposes. When it 
is dissolved in water, a considerable lowering of tempera- 
ture is caused. 

Ammonium Sulphide, (NH 4 ) 2 S. — This compound is 
extensively used in chemical analysis for the purpose of 
precipitating those sulphides which are soluble in dilute 
hydrochloric acid (see p. 214 and p. 215). As will be 
remembered, in the usual method of analyzing a mixture 
of substances, the first step consists in adding hydrochloric 
acid to the solution. This precipitates silver, lead, and, 
under certain conditions, mercury. This precipitate 
having been filtered off, hydrogen sulphide is passed 
through the filtrate, when those metals are precipitated 



~ -ii;>f s-'yliiifs :r: -.v. 5 ;"_7. .:1s in IiLv.:s ". ; .".: ; ;'_! ;:■-.: .v:i:.. 
The pr: j is filtered rf. and ammonium sulphide 

:ii: il:: ::: :1s ~:: lis - ?-: - :~_:.:rs Are 

soluble in dilate hydrochloric acid are thrown down. 
Am.:: these re iron, cobalt, nickel, manganese, etc 
other soluble sulphide migkr be use :: but the advan- 
tage of ammonium sulphide is that :: is volatile, and 
hen evaporating the slution and heath:. n be 

g : rid of after it has ebt bij 5 Another us 

which it is put in analysis is for the purpose of dissolving 
the sulphides of tin. arsenic, and antimony, which are 
precipitated by hydrogen sulphide, and thus separating 
these from the other sulphides of the group. This solu- 
ricn Lfy-fr. .".?"."■_ "-::: :n: : ■ :i ;: riis r.:l: :irs: : :_ 5 :".:? 
of sulpho-aeids. as has been repeatedly explained. 

Ammonium sulphide is made by passing hydrogen sul- 
phide into an aqueous solution of ammonia. If the g - 
is passed until the solution is saturated, the product is the 
l:-. : .::5 :iyi::A: 

xe 5 - eyg = NB^HS 

mbf half this quantity : :' _ - ^ - 

:5 :". "..>■' 

- ^= - ■'-"- > 

~:: ji—r'-rST :: - :: : " 

quantirj if a c lution of ammonia into two equal parrs : 
gainiate one half, thus forming the hydrosnlphide, and 
add the other half, when this rea :ion takes pi 

HXH 4 S - SH, = pro, - 

The product is a colorless liquid of a disagreeable odor. 
R s:-on changes color, becoming yellow, and after a time 
low der ; v .: is f rmed in the vessel in which it is con- 
tained. This changs uf color is i n of the 
ren of the air. S of the sulphic jiposed 
into ammot: . and sulphur, tt 3 

HH, >- = SHB.+fl/] - - 



SODIl \\l-L\f\fONIUM PHOSPHATE. 4*5 

The sulphur set free in this way combines with the un- 
decomposed ammonium sulphide, forming the compounds 
(NH 4 ),S„ (NHJ.S,, (NH 4 ) 2 S 4 , and (NHJ&. When as 
much sulphur has been set free as is required to form the 
pentasulphide, further decomposition by the oxygen of the 
air causes a deposit of sulphur. Therefore, in bottles 
containing ammonium sulphide which are allowed to stand 
for a long time a deposit of sulphur is always found. A 
solution containing the polysulphides is called yellow 
ammonium sulphide. It is this which is used for the pur- 
pose of dissolving the sulphides of arsenic, antimony, and 
tin in analytical operations. 

As stated above, a solution of ammonium hydrosulphide, 
11X11 + S, is made by passing hydrogen sulphide into a solu- 
tion of ammonia until no more is taken up. 

Ammonium Nitrate. NH 4 N0 5 , is obtained in crystals, 
which are easily soluble in water. It is of use chiefly in 
the preparation of nitrous oxide. When heated suddenly 
to a high temperature it is decomposed rapidly into 
nitrogen, water, and nitric oxide: 

BNH 4 N0, = X 2 + 2NO + 4H 2 0. 

This decomposition may take place in the preparation of 
nitrous oxide if in the last stages of the operation the heat 
is raised too high, and explosions may be caused in this 
way. When dissolved in water a marked lowering of tem- 
perature takes place. 

Sodium-ammonium Phosphate. HXaXH 4 P0 4 . — This 
salt is known as microcosmic salt, and is much used in the 
laboratory in blowpipe work. It is contained in guano 
and in decomposed urine. It is easily made by mixing 
solutions of di-sodium phosphate and ammonium chloride, 
and allowing to crystallize. In crystallized form it con- 
tains four molecules of water. IIXaXH 4 P0 4 -f 4H a O. The 
changes which the anhydrous salt undergoes when heated 
were described on page -AS;!. "When the crystallized salt 
is heated, the water of crystallization is first given off. 
The value of the salt in blowpipe work depends upon the 



486 COLLEGE CHEMISTRY. 

fact that at high temperatures the metaphosphate com- 
bines with metallic oxides, forming mixed phosphates, the 
reactions being like those which metaphosphoric acid 
undergoes with water: 

2HP0 3 + H 2 = H 4 P 2 7 ; 
HPO3 + H 2 = H 3 P0 4 ; 
2NaP0 3 + M 2 = Na 2 M 2 P 2 7 ; 
NaP0 8 + M 2 = NaM 2 P0 4 . 

Many of these double phosphates and pyrophosphates are 
colored, and furnish a means of detecting some of the 
metals. 

Reactions of the Members of the Sodium Group which 
are of Value in Chemical Analysis. — The chief difficulty 
experienced in chemical analysis is in distinguishing 
between similar elements. Sodium and potassium, for 
example, conduct themselves so much alike in so many 
respects that we might subject them to the influence of a 
number of reagents without being able to tell which one we 
are working with. For purposes of analysis, therefore, it 
is necessary to take advantage of differences between the 
elements, and the more striking the differences the better. 
Those reactions which give rise to the formation of insolu- 
ble compounds or precipitates are most frequently used in 
analysis. Very few salts of the members of the sodium 
group are insoluble, and the difficulty of distinguishing 
between these elements is increased by this fact. In 
ordinary analyses the elements of this group which are of 
most importance are potassium and sodium, the other 
elements of the group being but rarely met with. Am- 
monium compounds are easily distinguished from those of 
potassium and sodium by the fact that, when treated with 
caustic soda or potash, they give off ammonia, which is 
recognized by its characteristic odor. The chief reactions 
which are of value in distinguishing between potassium 
and sodium are the following : 

Chlorplatinic Acid, H 2 PtCl 4 , forms difficultly soluble 
salts with potassium and ammonium chlorides. These 



FLAME REACTIONS AND THE SPECTROSCOPE. 487 

are the chlorplatinates, K 2 PtCl 6 and (NH 4 ) a PtCl 6 . The 
corresponding salt of sodium is easily soluble. 

Perchloric Acid, HC10 4 , forms difficultly soluble potas- 
sium per chlorate, KC10 4 , when added to solutions of 
potassium salts. 

Fluosilicic Acid, H 2 SiF 6 , forms difficultly soluble salts 
with potassium and sodium, K 2 SiF 6 and Na 2 SiF 6 , but not 
with ammonium. 

Tartaric Acid, II 2 (C 4 H 4 6 ), forms a difficultly soluble 
potassium salt of the formula KH(C 4 H 4 6 ). The corre- 
sponding salt of sodium is easily soluble. The formation 
of mono-potassium tartrate takes place as represented in 
the equation: 

KCl + H 3 (C 4 H 4 6 ) = KH(C 4 H 4 6 ) + HC1. 

Normal or neutral potassium tartrate is soluble in water, 
so that, if the difficultly soluble acid tartrate is filtered 
off, and potassium carbonate added to it, it dissolves in 
consequence of the formation of the neutral salt, which 
takes place as represented in the equation 

2KH(C 4 H 4 6 ) + K 2 C0 3 = 2K,(C 4 H 4 6 ) + CO, + H,0. 

If, to the solution of the neutral salt, hydrochloric acid is 
added, the acid salt is again formed and precipitated : 

K,(C 4 H 4 6 ) + HC1 = KH(0 4 H 4 6 ) + KOI. 

Di-sodium Pyro-antimonate, Na 2 H 2 Sb 2 7 , is insoluble 
in cold water, and is formed when a solution of the corre- 
sponding potassium salt is added to a solution of a sodium 
salt. 

Flame Reactions and the Spectroscope. — When a clean 
piece of platinum wire is held for some time in the flame 
of the Bunsen burner, it then imparts no color to the 
flame. If now a small piece of sodium carbonate or any 
other salt of sodium is put on it, the flame is colored in- 



488 COLLEGE CHEMISTRY. 

tensely yellow. All sodium compounds have this power, 
and hence the chemist makes use of this fact for the pur- 
pose of detecting the presence of sodium. Similarly, 
potassium compounds color the flame violet ; lithium com- 
pounds color the flame red; rubidium and caesium produce 
colors similar to that of the potassium flame. While it is 
an easy matter to recognize potassium alone, or any one of 
the other metals alone, it is difficult to do so when they 
are together in the same compound. For example, when 
sodium and potassium are together, the intense yellow 
caused by the sodium completely masks the more delicate 
violet caused by the potassium, so that the latter cannot 
be seen by the unaided eye. In this particular case the 
difficulty can be got over by letting the light from the 
flame pass through a blue glass, or through a thin vessel 
of glass containing a solution of indigo. The yellow light 
is thus cut off, while the violet light passes through and 
can be recognized. A more general method for detecting 
the constituents of light is by means of a prism of glass. 
Lights of different colors, which are produced by ether 
waves of different lengths, are turned out of their course 
to different extents when passed through a prism, as is 
seen when white sunlight is passed through a prism. A 
narrow beam of white light passing in emerges as a band 
of various colors, called its spectrum. We thus see that 
white light is made up of lights of different colors; or, to 
speak in the language of physics, that motion of the light- 
ether which produces upon the eye the sensation of white 
light is made up of a number of motions, each of which 
alone produces upon the eye the sensation of a color. 
Similarly, we can determine what any light is composed 
of. Every light has its characteristic spectrum. The 
light given off from any solid heated to a white heat gives 
a continuous spectrum, like that of the sunlight. An in- 
candescent gaseous substance, on the other hand, gives a 
spectrum made up of separate bands of color, or a line 
spectrum. The light produced by burning sodium, or by 
introducing a sodium compound in a colorless flame, gives 



EXPERIMENTS WITH POTASSIUM SALTS. 4$9 

a spectrum consisting of ;i narrow yellow band. The 
spectrum of the potassium flame consists essentially of two 
bands, one red and one violet. Further, these bands 
always occupy definite positions relatively to one another, 
so that, in looking through a prism at the light caused by 
potassium and sodium, the yellow band of sodium is seen 
in its position, and the two potassium bands in their 
proper positions. There is therefore no difficulty in 
detecting these elements when present in the same sub- 
stance or in the presence of other elements which give 
characteristic spectra. 

The instrument used for the purpose of observing the 
spectra of different lights is called the spectroscope* It 
consists essentially of a prism and two telescopes. Through 
one of the telescopes the light to be examined is allowed 
to pass so as to strike the prism properly. The light 
emerges from the other side of the prism, and is observed 
through the other telescope, which is provided with lenses 
for the purpose of magnifying the spectrum. By means 
of a third telescope, an image of a scale is thrown upon 
the face of the prism from which the spectrum emerges, 
and is reflected thence into the obseiwing-tube, together 
with the spectrum, so that the position of the bands can 
be accurately determined. By means of the spectroscope, 
it is possible to detect the minutest quantities of some 
elements, and, since it was devised, several new elements 
have been discovered through its aid; as, for example, 
caesium, rubidium, thallium, indium, gallium, and others. 

EXPERIMENTS. 

Potassium Salts. 

Experiment 209.— In preparing potassium iodide from iodine 
and potassium hydroxide, proceed as follows : To 30 grams iodine 

* For an account of the spectroscope and its uses, the student 
should consult some work on physics. The principles involved in its 
construction and application are physical principles, and cannot 
properly be taken up in detail in a text-book of chemistry. 



49° COLLEGE CHEMISTRY. 

use 15 grams hydroxide. Dissolve the latter in 100 cc. water. 
Add half this solution to the iodine in a porcelain evaporating- 
dish. Now slowly add the rest of the liquid until the color dis- 
appears. Concentrate the liquid to a syrupy consistence, add 1 
gram finely-powdered charcoal, mix. and evaporate to dryness. 
The residue is then heated to redness in an iron vessel. After 
cooling extract with water. 

Experiment 210. — Potassium iodide can also be prepared by 
the following method: Bring together in a capsule 200 grams 
water, 10 grams iron filings, and 40 grams iodine ; mix, and heat 
gently. When the solution has become green, decant, filter, and 
wash. Now heat the liquid nearly to boiling, and gradually add 
a solution of 35 grams potassium carbonate in 100 grams water. 
Filter, wash, and evaporate. 

Experiment 211. — Dissolve 50 grams potassium carbonate in 
500 to 600 cc. water. Heat to boiling in an iron or a silver ves- 
sel, and gradually add the slaked lime obtained from 25 to 30 
grams of good quicklime. During the operation the mass should 
be stirred with an iron spatula. After the solution is cool, draw 
it off by means of a siphon into a bottle. This may be used in 
experiments in which caustic potash is required. 

Experiment 212. — Mix together 15 grams potassium nitrate 
and 2.5 grams powdered charcoal. Set fire to the mass. 

Experiment 213. — Treat a quantity of wood-ashes with water. 
Filter, and examine by means of red litmus paper. Evaporate 
to dryness. What evidence have you that the residue contains 
potassium carbonate ? 



Sodium Salts. 

Experiment 214. — Make a supersaturated solution of sodium 
sulphate by heating an excess of the salt with water at 33°. 
Filter the solution into small flasks and cork them. On remov- 
ing the corks and agitating the vessels, the salt will suddenly 
crystallize out. 

Experiment 215.— Make a saturated solution of common salt 
in ordinary ammonia-water (about 50 cc). Pass carbon dioxide 
into this solution until no more is absorbed, the delivery-tube 
being arranged as shown in Experiment 75. Filter off the pre- 
cipitate, and dry it by spreading it upon layers of filter-paper. 
Heat some of the salt when dry, and determine whether the gas 
given off is carbon dioxide or not. When gas is no longer given 



EXPERIMENTS WITH SODIUM SALTS. 49 1 

off by heating, let the tube cool and examine the residue. [Is it 
a carbonate ?] 

Experiment 210. — Make ammonium sulphide thus : Divide a 
given quantity of a solution of ammonia into two equal parts. 
Saturate one half by passing hydrogen sulphide through it, and 
then add the other half. 



CHAPTER XXVI. 

ELEMENTS OF FAMILY II, GROUP A: 

GLUCINUM.— MAGNESIUM.— CALCIUM.— STRONTIUM. 

—BARIUM [ERBIUM]. 

General. — The elements of this group fall into two sub- 
groups. Calcium, strontium, and barium are strikingly 
alike. They also have some points in common with the 
members of the potassium family, and at the same time 
are related in some degree to the metals of Family III, 
Group A, which are known as the earth metals. There- 
fore, calcium, barium, and strontium are generally called 
the metals of the alkaline earths. Glucinum and magnesium 
resemble the metals of the alkaline earths in some ways, 
but they also resemble the members of Group B, of the 
same family, which includes zinc and cadmium. On com- 
paring the group with the elements presented in the last 
chapter, some analogous facts are noticed. Arranging the 
five elements of the potassium group in the order of their 
atomic weights, and the elements of Family II, Group A, 
in the same way, we have this table : 



Li 


Na 


K 


Rb 


Cs 


7.03 


23.05 


39.15 


85.4 


133 


Gl 


Mg 


Ca 


Sr 


Ba 


9.1 


24.36 


40 


87.6 


137.4 



As regards the analogies between the elements in each 
group, the last three members of each group resemble one 
another more closely than they resemble the first two 
members of the group, while the first two members in each 
group also resemble each other closely. The natural 

492 



■ 



ELEMENTS OF FAMILY II, GROUP A. 493 

grouping according to the properties is into the sub- 
groups : 



Lithium, 
Sodium, 



b 
Potassium, 
and Kubidium, 

Caesium. 



n , . Calcium, 

Glucmum, _ _., ,. 

tv T . and Strontium, 

Magnesium, _, . ' 

Barium. 

The relations between the atomic weights of the elements 
of Family II, Group A, are similar to those of the elements 
of Family I, Group A. That of magnesium, 24.36, is 
nearly half the sum of those of glucmum, 9.1, and cal- 
cium, 40. We have 

= 24.55. 

So, also, that of strontium, 87.6, is approximately half 
the sum of those of calcium, 40, and barium, 137.4. 

40 + 137.4 = 88y _ 

/J 

In the calcium group the specific gravities increase in the 
order of the atomic weights : 

At. Wt. Sp. Gr. 

Calcium 40 1.57 

Strontium 87.6 2.5 

Barium 137.4 3.75 

All the elements of the group are bivalent. The general 
formulas of the principal compounds are as follows : 

MC1 2 , M(OH) 2 , M(NO s ) 2 , MS0 4 , M 3 (POJ 2 , MSi0 3 , etc. 

The chlorides, hydroxides, and nitrates are soluble in 
water. The sulphates decrease in solubility as the atomic 
weights increase. Glucinum sulphate, G1S0 4 , is soluble 



494 COLLEGE CHEMISTRY. 

in its own weight of water; magnesium sulphate, MgS0 4 , 
is soluble in about three times its weight of water; calcium 
sulphate, CaS0 4 , dissolves in 400 parts; strontium sul- 
phate, SrS0 4 , in about 8000 parts; and barium sulphate, 
BaS0 4 , in about 400,000 parts of water. Barium sul- 
phate, as will be seen, is practically insoluble in water. 
The normal carbonates of all except glucinum are insoluble 
in water. The solubility of the hydroxides increases as 
the atomic weight increases. Glucinum hydroxide is in- 
soluble; magnesium hydroxide is but slightly soluble. 
One hundred parts of water at the ordinary temperature 
dissolve 0.1368 part of calcium hydroxide, 2 parts of 
strontium hydroxide, and 3.5 parts of barium hydroxide. 
The solubility of strontium and barium hydroxides is, 
however, much increased at higher temperatures. 



CALCIUM SUB-GROUP. 

This sub-group, as has been stated, consists of the three 
similar elements, calcium, strontium, and barium. Of 
these calcium occurs most abundantly in nature. Barium 
and strontium frequently accompany each other, and both 
are found in some localities in company with calcium. 
They are much less abundant in nature than calcium. 

Calcium, Ca (At. Wt. 40). 

Occurrence. — Calcium is found in nature in enormous 
quantities, chiefly in the form of the carbonate, CaC0 3 , 
as limestone, marble, and chalk. It also occurs in the 
form of the sulphate, CaS0 4 , as gypsum; of the phos- 
phate, Ca 3 (P0 4 ) 2 , as phosphorite and apatite; of the 
fluoride, CaF 2 , as fluor-spar. It is found in solution in 
most natural waters either as the carbonate or sulphate; 
and in the organs of plants and animals. Bones contain 
a large proportion of calcium phosphate; egg-shells and 
coral contain calcium carbonate. 



CALCIUM. 495 

Preparation. — Calcium is made: (1) by heating together 
anhydrous calcium iodide and metallic sodium; (2) by 
heating calcium oxide and carbon together in an electric 
furnace in a current of hydrogen. In the latter case a 
mixture of calcium and the hydride is obtained, but the 
metal can be separated from this without difficulty. 

Properties. — Calcium is a silver-white metal. It is soft 
and can be cut with a knife like sodium. It is one of the 
most active elements known, and combines with all the 
other elements except argon. It combines with hydrogen 
to form the hydride CaH 2 , and burns brilliantly in the air, 
forming compounds with both oxygen and nitrogen. It 
decomposes water easily and becomes incandescent if in 
the form of powder. 

Calcium Chloride, CaCl 2 . — This salt is found in nature 
in combination with other chlorides, particularly in the 
mineral tachydrite, which occurs in the salt deposits at 
Stassfurt, and has the composition represented by the 
formula CaCl 2 .MgCl 2 + 12H 2 0. It is also found in solu- 
tion in sea-water. It is obtained as a by-product in the 
preparation of ammonia from ammonium chloride and 
lime; in the preparation of potassium chlorate from cal- 
cium chlorate and potassium chloride (see p. 464); and 
in the ammonia-soda process. It is made by dissolving 
calcium carbonate in hydrochloric acid, as in the prepara- 
tion of carbon dioxide. From very concentrated solutions 
it crystallizes with six molecules of water, CaCl 2 -(- 6H 2 0. 
When these crystals are exposed to the air they soon 
deliquesce. When a solution of calcium chloride is evap- 
orated, and care is taken to keep the temperature below 
200°, it solidifies, forming a porous mass which has the 
composition represented by the formula CaCl 2 -\- 2H 2 0. 
This is much used in laboratories as a drying agent, as it 
absorbs water with great ease. If this salt is heated above 
200° it loses all its water, and the dehydrated chloride 
melts, forming fused calcium chloride. This is also much 
used on account of its drying power. Gases are passed 
through tubes filled with granulated calcium chloride for 



49 6 COLLEGE CHEMISTRY. 

the purpose of drying them, and the salt is also placed in 
vessels in which it is necessary that the air should be dry, 
as in balance-cases, desiccators, etc. 

Calcium chloride forms crystallized compounds with 
ammonia and with alcohol, as well as with water. It is 
obvious from this that calcium chloride cannot be used for 
the purpose of drying ammonia gas. When the compounds 
with ammonia and with alcohol are heated they break 
down, yielding ammonia and alcohol respectively, as the 
compound with water gives up the latter. 

Calcium Fluoride, CaF 2 .— This compound occurs in 
large quantities in nature as the mineral fluor-spar. It 
occurs beautifully crystallized in cubes, and is insoluble in 
water. It is the source of fluorine compounds in general, 
and is used in metallurgical operations for the reason that 
it melts readily and does not act upon other substances 
easily. It therefore simply serves as a liquid medium in 
which reactions take place at high temperatures. A sub- 
stance that' acts in this way and is used for this purpose 
is called a flux. The name fluor-spar has its origin in 
this use of the substance. 

Calcium Oxide, CaO. — This important compound is 
commonly called lime, or, to distinguish it from the 
hydroxide or slaked lime, it is called quick-lime. It is 
made in large quantity by heating calcium carbonate in 
appropriately constructed furnaces, known as lime-kilns. 
Pure lime is made by decomposing some pure form of 
calcium carbonate, as marble or calc-spar. 

Lime is a white, amorphous, infusible substance. When 
heated in the flame of the compound blowpipe it gives an 
intense light, as any other infusible substance would do 
under the same circumstances. When exposed to the air 
it attracts moisture and carbon dioxide, and is converted 
into the carbonate. It must therefore be protected from 
the air, Lime which has been converted into the car- 
bonate by exposure to the air is said to be air-slaked. 

Calcium Hydroxide, Ca(OH) 2 . — When calcium oxide or 
quick-lime is treated with water it becomes hot and 



CALCIUM HYDROXIDE. 497 

crumbles to a fine powder. The substance formed in this 
operation is somewhat soluble in water, the solution being 
known as lime-water. The chemical change that takes 
place when lime is treated with water has been explained. 
It consists in the formation of a compound of the formula 
Ca(OH) 2 , known as slaked lime; and the operation is 
known as slaking. The action is of the same kind as that 
with which we have so frequently had to deal in the trans- 
formation of oxides into the corresponding hydroxides. 
Thus when potassium oxide is treated with water it is 
changed to the hydroxide, with a marked evolution of 
heat, the reaction being represented in this way: 

K 2 + H 2 = fgg. 

So, too, when sulphur trioxide is brought in contact with 
water it appears to form the hydroxide, normal sulphuric 
acid: 

fOH 

OH 

OH 

OH* 

OH 
[OH 



( = 
S 1 = + 3H 2 = S 1 
-0 



The action in the case of calcium oxide is represented in 
a similar way : 

Ca = + H 2 = Ca<°|. 

The hydroxide is a fine white powder. At red heat it 
loses water and is reconverted into the oxide : 

Ca<^ = Ca = + H 2 0. 

When lime-water is exposed to the air it becomes covered 
with a crust of calcium carbonate, and finally all the cal- 
cium is precipitated as calcium carbonate. A solution of 
calcium hydroxide affords a convenient means of detecting 



498 COLLEGE CHEMISTRY. 

the presence of carbon dioxide, as has been shown in deal- 
ing with this gas. The solution has an alkaline reaction, 
and acts in many respects like solutions of the hydroxides 
of potassium and sodium. 

Applications. — Lime is extensively used in the arts, 
generally in the form of the hydroxide. As we have seen, 
it is used in the preparation of ammonia and the caustic 
alkalies, potassium and sodium hydroxides; and of bleach- 
ing-powder and potassium chlorate. It is further used in 
large quantity for the purpose of removing the hair from 
hides in the process of tanning; in decomposing fats for 
the purpose of making stearin for candles; for purifying 
gas ; and especially in the preparation of mortar. 

Bleaching-powder. — The preparation of bleaching- 
powder was referred to under Chlorine (which see). The 
main reaction involved is that represented in the equation 

20a(OH) 2 + 4C1 = Ca(C10) 2 + CaCl 2 + 2H 2 0. 



Y 

Bleaching-po wder, 



The compound is commonly called "chloride of lime." 
Assuming that the reaction takes place. in the same way a:s 
that of chlorine on caustic potash, the product is a mix- 
ture of calcium hypochlorite, Ca(C10) 2 , and calcium 
chloride, for it is held that the reaction with potassium 
hydroxide takes place as represented in this equation : 

2KOH + 2C1 = KCIO + KC1 + H 2 0. 

An objection to the view that calcium chloride is present 
as such in bleaching-powder is found in the fact that the 
substance is not deliquescent, as it should be if calcium 
chloride were present. This has led to the suggestion that 
bleaching-powder in the dry form is not a mixture of two 
compounds, as represented above, but that it is rather one 

compound of the formula Ca -j ~p, or CaOCl 2 . A com- 
pound of this formula would plainly have the same com- 
position as a mixture of calcium hypochlorite and calcium 



BLEACHING-POWDER. 499 

chloride in the proportion of their molecular weights. 
For we have 

Ca(C10) 2 + CaCl 2 = 2CaOCl 2 . 

The point is a difficult one to decide, but at present the 
evidence appears to be rather in favor of the view that 
bleaching-powder in the dry form is a single compound of 
the constitution represented by the last formula given. 
When treated with water, however, it conducts itself like 
a mixture of the hypochlorite and chloride. This is 
entirely in accordance with the theory of electrolytic dis- 
sociation, as the ions would be the same in both cases: 

CaCl 2 = Ca + 2C1; 
Ca(OCl) 2 = Ca + 20C1. 
2Ca<^ 1 C1 = 2Ca + 20C1 + 2~cT. 

Bleaching-powder is a white powder which has the odor 
of hypochlorous acid. It is soluble in about twenty parts 
of water, though the commercial product always leaves a 
slight residue, which consists mainly of calcium hydroxide. 
When treated with an acid, as sulphuric or hydrochloric 
acid, it gives up all its chlorine. Thus, with hydrochloric 
acid the reaction takes place as represented in these equa- 
tions : 

Ca(C10) 2 + 2HC1 = CaCl 2 + 2HC10; 
2H01 + 2HC10 = 2H 2 + 2C1 2 . 

With sulphuric acid the action also probably takes place 
in two stages. The acid acts upon the hypochlorite, 
setting hypochlorous acid free; and upon the chloride, 
setting hydrochloric acid free. The hydrochloric and 
hypochlorous acids then react with each other as repre- 
sented above : 

Ca(C10) 2 + H 2 S0 4 = CaS0 4 + 2HC10; 
CaCl 2 -f H 2 S0 4 = CaS0 4 + 2HC1; 
2HC1 + 2HC10 = 2H 2 + 2C1 2 . 



500 COLLEGE CHEMISTRY. 

"When exposed to the action of carbon dioxide hypo- 
chlorous acid is liberated. Hence, when it is allowed to 
lie in the air this decomposition takes place slowly. The 
hypochlorous acid acts further upon the calcium chloride, 
liberating chlorine: 

CaCl 2 + 2HOC1 + C0 2 = CaC0 3 + H 2 + 2C1 2 . 

It may be, however, that the action takes place between 
carbon dioxide and the compound CaOCl 2 , thus: 

CaOCl 2 + C0 2 = CaC0 3 + Cl 2 . 

In any case, the fact remains that carbon dioxide sets the 
chlorine free from bleaching-powder. 

A solution of bleaching-powder alone is not capable of 
bleaching except very slowly. If, however, something is 
added which has the power to decompose it, bleaching 
takes place, the action being due to the presence of hypo- 
chlorous acid and chloriue. As is clear from what was 
said above the passage of carbon dioxide through the 
solution or the addition of an acid would cause it to bleach. 
So, too, certain salts produce a similar effect. The ex- 
planation of this is the instability of the hypochlorites 
formed by the salts added. When a concentrated solution 
of bleaching-powder is heated it gives off oxygen, and the 
salt is converted into the chloride. In dilute solution, 
however, the hypochlorite is converted into chlorate and 
chloride : 



3Ca(C10) 2 = Ca(C10 3 ) 2 + 2CaCl 



This fact is taken advantage of, as has been shown, for 
the purpose of making calcium chlorate, and from this 
potassium chlorate (see p. 464). In contact with certain 
oxides, as copper oxide, ferric oxide, and with hydroxides, 
as cobalt and nickel hydroxides, a solution of bleaching- 
powder readily gives up oxygen when heated. 

Applications. — The chief application of bleaching- 
powder is, as its name implies, for bleaching. It is also 
used as a disinfectant, and as an antiseptic, that is, for the 



CALCIUM CARBONATE. S 01 

purpose of destroying disease germs, and of preventing 
decomposition of organic substances. 

Calcium Carbonate, CaC0 3 .— This salt occurs in im- 
mense quantities in nature in the well-known forms lime- 
stone, calc-spar, marble, and chalk. The variety of calc- 
spar found in Iceland, and known as Iceland spar, is particu- 
larly pure calcium carbonate. It crystallizes in a number 
of different forms, the most common being in rhoinbo- 
hedrons, as seen in ordinary calc-spar. A second variety 
of crystallized calcium carbonate is aragonite. This is 
found in nature crystallized in rhombic prisms, and in 
forms derived from this. When heated, aragonite falls to 
pieces, the particles being small crystals of the form char- 
acteristic of calc-spar. This is a case of dimorphism 
similar to that presented by sulphur, which, it will be 
remembered, crystallizes in two forms, the rhombic and 
monoclinic, the latter of which passes into the former 
spontaneously. These forms are produced artificially very 
readily. When calcium carbonate is precipitated from a 
solution of a calcium salt by adding a soluble carbonate at 
ordinary temperatures the precipitate is made up of micro- 
scopic crystals which have the same form as calc-spar. If, 
however, the solution from which the carbonate is precipi- 
tated is hot, the salt consists of microscopic crystals of the 
form of aragonite. 

The most abundant form of calcium carbonate is lime- 
stone, of which many great mountain-ranges are largely 
made up. This is a compact form of the compound, that 
has a gray color, and frequently consists of minute 
crystals. It is always more or less impure, containing clay 
and other substances. Limestone mixed with a consider- 
able proportion of clay is called marl. Many natural waters 
contain calcium carbonate in solution — probably in the 
form of the acid carbonate. When such a water evaporates 
the carbonate is again deposited. It happens in some 
places that a water charged with the carbonate works its 
way slowly through the earth and drops from the top of a 
cave. Under these circumstances there is a gradual 



5o* COLLEGE CHEMISTRY. 

deposit of the salt which remains suspended. Such hang- 
ing formations of the carbonate are known as stalactites. 
At the same time that part of the liquid which falls to the 
bottom of the cave forms a projecting mass below the 
stalactite. Such projecting masses are called stalagmites. 
The formation of stalactites takes place in much the same 
way as that of icicles. 

Much of the calcium carbonate found in nature has its 
origin in the remains of animals, and fossils are yery 
abundant in it. Chalk consists almost exclusively of the 
shells of microscopic animals. 

When carbon dioxide is passed into a solution of calcium 
hydroxide, the carbonate is precipitated; and, if the cur- 
rent of gas is continued long enough, the carbonate is 
redissolved. It appears, therefore, that calcium carbonate 
is soluble in water that contains carbonic acid. It is 
probable that the cause of this is to be found in the forma- 
tion of an acid carbonate. Natural waters which come in 
contact with limestone gradually take up more or less of 
the carbonate, with the aid of the carbon dioxide of the 
air, and when such a water is boiled, the carbonate is 
thrown down. A water containing calcium carbonate in 
solution is called a hard tvaterj and as this kind of hard- 
ness is easily removed by boiling, it is called temporary 
hardness in order to distinguish it from a kind which is 
not removed by boiling, and is therefore called permanent 
hardness. Temporary hardness is further removed by 
adding lime to the water, when normal carbonate is 
formed, which is at once precipitated. 

The decomposition of calcium carbonate by heat, lead- 
ing to the formation of lime, or calcium oxide, and carbon 
dioxide, was referred to on p. 450. 

Applications. — Calcium carbonate is used, in the arts, 
for a great many purposes, as in the manufacture of glass; 
as a flux (see p. 496) in many important metallurgical 
operations, as in the reduction of iron from its ores; in 
the preparation of lime for mortar; etc. As is well known, 
further, marble and some of the varieties of limestone are 



CALCIUM SULPHATE. 503 

extensively used in building; and large quantities of chalk 
are also used. 

Calcium Sulphate, CaS0 4 . — This compound is very 
abundant in nature. The principal natural variety is 
gypsum, which occurs in crystals containing two molecules 
of water, CaS0 4 + 2H 2 0, This is perhaps derived directly 
from the normal acid S(OH) 6 , having the constitution 

represented by the formula (HO) 4 S<^>Ca. The salt of 

the formula CaS0 4 also occurs in nature, and is called 
anhydrite. A granular form of gypsum is called alabaster. 
Calcium sulphate is difficultly soluble in hot and cold 
water, but its solubility is markedly increased by the pres- 
ence of certain other salts; as, for example, sodium 
chloride. It is comparatively easily soluble in hydrochloric 
acid and in nitric acid. When heated to 100°, or a little 
above, it loses nearly all its water and forms a powder 
known as plaster of Paris, which has the power of taking 
up water and forming a solid substance. This process of 
solidification is known as "setting." Plaster of Paris is 
very largely used in making casts, on account of its power 
to harden after having been made into a paste with water. 
The hardening is a chemical process, and is caused by the 
combination of water with the salt to form the crystallized 
variety : 

CaS0 4 + 2H 2 = (HO) 4 S<£>Ca. 

When heated to 200°, and above, all the water is given 
oif from gypsum, and the product now combines with 
water only very slowly, and is of no value for making 
casts. In general, the higher the temperature to which 
the gypsum is heated, the greater the difficulty with which 
the product combines with water. 

Many natural waters contain gypsum in solution. Such 
waters act in some respects like those which contain cal- 
cium carbonate. With soap, for example, they form 
insoluble compounds. They are called hard waters. 
This kind of hardness is not removed by boiling, and it is 



504 COLLEGE CHEMISTRY. 

therefore called permanent hardness. Magnesium sulphate 
acts in the same way, producing permanent hardness. 

When calcium sulphate is treated with a solution of a 
soluble carbonate, it is decomposed, forming the carbonate 
as represented in the equation 

OaS0 4 -f Na 2 C0 3 = Na 2 S0 4 + Ca00 3 . 

This change is effected by allowing the two to stand in 
contact at the ordinary temperature. 

Applications. — Besides being used for making casts, 
calcined gypsum is used in surgery for making plaster-of- 
Paris bandages, and as a fertilizer. Its action as a fertili- 
zer is believed by some to be due to the fact that it has 
the power to hold ammonia and ammonium carbonate in 
combination, and thus to make them available for the 
plants. It has recently been shown that it in some way 
facilitates the process of nitrification, and perhaps it is in 
consequence of this that it facilitates plant-growth. 

Calcium Phosphates. — There are three phosphates of 
calcium: (1) The normal phosphate, Ca 3 (P0 4 ) 2 ; (2) the 
secondary phosphate, CaHP0 4 ; and (3) the primary phos- 
phate, CaH 4 (POJ 2 . 

(1) Normal calcium phosphate, Ca 3 (P0 4 ) 2 , is derived from 
phosphoric acid by the substitution of calcium for all the 
hydrogen. It is found in nature in large quantity as 
phosphorite, and in combination with calcium fluoride or 
chloride as apatite. It is, further, the chief inorganic 
constituent of bones, forming 85 per cent of bone-ash, and 
is contained in the excrement of animals, as in guano, etc. 
It is found everywhere in the soil, and is taken up by the 
plants for whose development it is essential. That it is 
also essential to the life of animals is obvious from the fact 
that the bones consist so largely of it. The phosphate 
needed for the building up of bones is taken into the 
system with the food. From these statements, it is clear 
that calcium phosphate is of fundamental importance, and 
that a fertile soil must either contain this salt or something 
from which it can be formed. Now, when a crop is raised 



CALCIUM PHOSPHATES. 5° 5 

on a given area, a certain amount of the phosphate con- 
fcained in it is withdrawn. If the plants were allowed to 
decay where they grow, the phosphate would be returned 
and the soil would continue fertile; but in cultivated lands 
this is not the case. The crops are removed, and with 
them the calcium phosphate contained in them, and the 
soil therefore becomes exhausted. If the substances 
removed are used as food, some of the phosphate is found 
in the excrement of the animals; and, if this excrement 
is put on the soil, it is again rendered fertile. There are, 
however, other sources of calcium phosphate, and some of 
these are utilized extensively in the preparation of artificial 
fertilizers. The natural form of the phosphate, as that in 
bone-ash, in phosphorite, and in guano, is mainly 'the 
normal or neutral phosphate. This is insoluble in water, 
and is therefore taken up by the plants with difficulty. 
To make it quickly available, it must be converted into a 
soluble phosphate. This is done by treating it with sul- 
phuric acid in order to effect the reaction represented in 
this equation: 

Ca 3 (P0 4 ) 2 -j- 2H 2 S0 4 = CaH 4 (P0 4 ) 2 + 2CaSO,. 

The primary phosphate thus formed is soluble in water, 
and is of great value as a fertilizer. The mixture of the 
soluble phosphate and of calcium sulphate is known as 
" superphosphate of lime." The sulphate, as we have 
seen, is also of value as a fertilizer. The value of super- 
phosphates depends chiefly upon the amount of soluble 
phosphate contained in them ; and in dealing with them it 
is customary to state how much "soluble" and how much 
"insoluble phosphoric acid" they contain. When a 
superphosphate is allowed to stand for a time, some of the 
soluble primary phosphate is converted into insoluble 
phosphates by contact with basic hydroxides and water. 
This is known as the 'process of " reversion," and that part 
of the phosphoric acid which is contained in the insoluble 
phosphate is spoken of as "reverted phosphoric acid." 
Normal calcium phosphate, as has been stated, is in- 



506 COLLEGE CHEMISTRY. 

soluble in water, and is formed when a soluble normal 
phosphate is added to a solution of a calcium salt. It is 
also formed when di-sodium phosphate and ammonia are 
added to a solution of a calcium salt, thus : 

2HNa 2 P0 4 + 3CaCl 2 + 2NH 3 = Ca 3 (P0 4 ) 2 -f 4NaCl + 2NH 4 C1. 

Normal or tertiary calcium phosphate is soluble in 
hydrochloric acid and in nitric acid, in consequence of the 
formation of calcium chloride, or nitrate, and the primary 
phosphate. If ammonia is added to this solution, the 
tertiary phosphate is again precipitated, as represented 
below : 

Ca 3 (P0 4 ) 2 + 4HC1 = 2CaCl 2 + H 4 Ca(P0 4 ) 2 ; 
20aCl 2 + H 4 Ca(P0 4 ) 2 -f 4NH 3 = Ca 3 (P0 4 ) 2 + 4NH 4 01. 

(2) Secondary calcium phosphate, CaHP0 4 , is formed, 
as above described, when a solution of a calcium salt is 
treated with secondary sodium phosphate. 

(3) Primary calcium phosphate, H 4 Ca(P0 4 ) 2 , is com- 
monly called the acid phosphate of calcium. It is formed 
when ordinary insoluble calcium phosphate is treated with 
concentrated sulphuric acid, and is contained in the 
so-called superphosphates. It is also formed by treating 
the neutral phosphate with phosphoric acid and with 
hydrochloric acid. When treated with but little water, it 
is converted into the secondary salt and free acid : 

' H 4 Ca(P0 4 ) 2 = HCaP0 4 + H 3 P0 4 . 

Calcium Silicate, CaSi0 3 , occurs in nature as the 
mineral wollastonite, and, in combination with other sili- 
cates, in a large number of minerals, as garnet, mica, the 
zeolites, etc. It is formed when a solution of sodium sili- 
cate is added to a solution of calcium chloride, and when 
a mixture of calcium carbonate and quartz is heated to a 
high temperature. 

Glass. — Ordinary glass is a silicate of calcium and 
sodium made by melting together sand (silicon dioxide, 
Si0 2 ) with lime and sodium carbonate or soda. Instead 



GLASS. 5°7 

of calcium carbonate, lead oxide may be used ; and instead 
of sodium carbonate, potassium carbonate. The proper- 
ties of the glass are dependent upon the materials used in 
its manufacture. 

Ordinary window glass is a sodium-calcium glass. The 
purer the calcium carbonate and silica, the better the 
quality of the glass. This glass is comparatively easily 
acted upon by chemical substances, and is therefore not 
adapted to the preparation of vessels which are to be used 
to hold acids and other chemically active substances. It 
answers, however, very well for windows. The difference 
between ordinary window glass and plate glass is essentially 
that the former is blown and then cut into pieces, while 
the latter, when in the molten condition, is run into flat 
moulds and there allowed to solidify. 

Bohemian glass is made with potassium carbonate. If 
pure carbonate is used, as well as pure calcium carbonate 
and silica, a very beautiful glass is the result. It is char- 
acterized by great hardness, by its difficult fusibility, and 
by its resistance to the action of chemical substances. It 
is particularly well adapted to the manufacture of vessels 
and tubes for use in chemical laboratories. 

Flint-glass is made by melting together lead oxide, 
potassium carbonate, and silicon dioxide. It is character- 
ized by its power to refract light, its high specific gravity, 
its low melting-point, and the ease with which it is acted 
upon by reagents. Owing to its high refractive power, it 
is largely used in the manufacture of lenses for optical 
instruments. 

Strass is a variety of lead-glass that is particularly rich 
in lead. Its refracting power is so great that it is used in 
the manufacture of artificial gems. 

Colors are given to glass by putting in the fused mass 
small quantities of various substances. Thus, a cobalt 
compound makes glass blue; copper or chromium makes 
it green; one of the oxides of copper makes it red; uranium 
gives it a yellow color; etc. The most common variety of 
glass is that used in the manufacture of ordinary bottles. 



508 COLLEGE CHEMISTRY. 

It is generally green to black, and sometimes brown. In 
its manufacture, impure materials are used, chiefly ordi- 
nary sand, limestone, sodium sulphate, common salt, clay, 
etc. 

When glass is suddenly cooled, it is very brittle and 
breaks into small pieces when scratched or slightly broken 
in any way. This is shown by the so-called Prince 
Rupert's drops, which are made by dropping glass, in the 
molten condition, into water. When the end of such a 
drop is broken off, the entire mass is completely shattered 
into minute pieces. It is clear from this that, in the 
manufacture of glass objects, care must be taken not to 
cool them suddenly. In fact they are cooled very slowly, 
the process being known as annealing. For this purpose 
they are placed in furnaces the temperature of which is 
but little below that of fusion, and they are kept there for 
some time, the heat being gradually lowered. If red-hot 
glass is introduced into heated oil or paraffin, and allowed 
to cool very slowly, it is found to be extremely hard and 
elastic. The glass of De la Bastie is made in this way. 
Vessels made of it can be thrown about upon hard objects 
without breaking, but sometimes a slight scratch will cause 
the glass to fly in pieces as the Rupert's drops do. 

Mortar. — Mortar is made of slaked lime and sand. 
When this mixture is exposed to the air, calcium carbonate 
is slowly formed and the mass becomes extremely hard. 
The water contained in the mortar soon passes off, but 
nevertheless freshly plastered rooms remain moist for a 
considerable time. This is due to the fact that a reac- 
tion takes place between the carbon dioxide and calcium 
hydroxide by which calcium carbonate and water are 
formed, 

Ca(OH) 2 + C0 2 = CaC0 3 + H 2 0, 

and it is the water thus liberated that keeps the air 
moist. The complete conversion of the lime into car- 
bonate requires a long time, because the carbonate which 



MORTAR-CALCIUM SULPHIDE. 509 

is formed on the surface protects, to some extent, the lime 
in the interior. 

It is generally regarded as unhealthy to live in rooms 
with freshly plastered walls, because the air is constantly 
kept moist in consequence of the reaction above mentioned. 
It is, however, difficult to see why the presence of a little 
extra moisture in the air should be unhealthy; and, if 
there is any danger from freshly plastered walls, it seems 
probable that the cause must ba sought for elsewhere. It 
is possible that the constant presence of moisture in the 
pores of the wall interferes with the important process of 
diffusion, and that therefore when the room is closed this 
natural method of ventilation cannot come into play. 

When limestones which contain magnesium carbonate 
and aluminium silicate in considerable quantities are 
heated for the preparation of lime, the product does not 
act with water as calcium oxide does, and this lime is not 
adapted to the preparation of ordinary mortar. On the 
other hand, it gradually becomes solid, in contact with 
water, for reasons which are not known. Such substances 
are known as cements, or hydraulic cements. 

Calcium Sulphide, CaS, is formed by heating calcium 
sulphate with charcoal. It is converted by water into the 
hydrosulphide and the hydroxide: 

2CaS + 2H 2 = Ca(SH) 2 + Ca(OH) 2 . 

The hydrosulphide is easily soluble in water, while the 
hydroxide is difficultly soluble. If the water is saturated 
with the hydroxide, it does not act upon the sulphide. It 
is remarkable on account of the fact that it is phosphores- 
cent. After having been exposed to sunlight, it continues 
to give light for some time afterward. This and the 
similar compound, barium sulphide, are now used in the 
preparation of luminous objects, such as match-boxes, 
clock-faces, plates for house-numbers, etc. 

Calcium Nitride, Ca 3 N 2 , a bronze-colored mass, is 
formed by heating calcium in nitrogen. It is decomposed 



5*° COLLEGE CHEMISTRY. 

by water, with the formation of calcium hydroxide and 
ammonia. 

Calcium Carbide, C 2 0a, is formed by heating lime and 
carbon together in an electric furnace, when the reaction 
represented by the following equation takes place : 

CaO + 30 = C 2 Ca + CO. 

It is manufactured on the large scale, the form of carbon 
used being coke. The carbide forms a crystalline mass. 
That of average quality has a reddish color. That of bad 
quality has a grayish or black color. When pure it is 
transparent and white. When treated with water it yields 
acetylene, C 2 H 2 (which see). 

Strontium, Sr (At. Wt. 87.6). 

Occurrence and Preparation. — Strontium occurs in 
nature in the form of the sulphate, SrS0 4 , as celestite, 
and in the form of the carbonate, SrC0 3 , as strontianite. 
The latter is found in large quantities in Westphalia. 
The element is isolated by the action of an electric current 
on the molten chloride. 

Properties. — It is similar to calcium. It is oxidized by 
contact with the air, and decomposes water rapidly with 
evolution of hydrogen, which does not, however, take fire 
spontaneously. 

Compounds of Strontium. — The compounds of strontium 
are similar to those of calcium. Its chloride has not the 
same power to combine with water that calcium chloride 
has, though it deliquesces when left in contact with the 
air. The oxide- is not easily made by decomposition of the 
carbonate by heat, as the carbonate is much more stable 
than that of calcium. It is, however, prepared without 
difficulty by heating the nitrate. When brought in con- 
tact with water, the oxide forms the hydroxide, which is 
analogous to calcium hydroxide. It is more easily soluble 
in water than the latter. 

Strontium nitrate, Sr(N0 3 ) 2 , is made in considerable 



BARIUM. 5 11 

quantity for the purpose of preparing a mixture which, 
when burned, gives a red light (red-fire, Bengal-fire). It 
is easily made by dissolving strontianite or strontium car- 
bonate in nitric acid. 

Strontium sulphate, SrS0 4 , occurs in nature in beauti- 
ful crystals as the mineral celestite. It is formed when a 
soluble sulphate is added to a solution of a strontium salt. 
Its solubility lies between that of calcium sulphate and 
that of barium sulphate. 

Barium, Ba (At. Wt. 137.4). 

Occurrence and Preparation. — Barium occurs in nature 
in the same forms of combination as strontium, viz., as 
the carbonate, BaC0 3 , in witherite; and as the sulphate, 
BaS0 4 , in barite or heavy spar. It is prepared by electrol- 
ysis of the molten chloride. 

Properties. — It closely resembles calcium and strontium, 
being a yellow metal, which is oxidized by contact with 
the air and readily decomposes water at the ordinary tem- 
perature. 

Barium Chloride, BaCl 2 -f 2H 2 0, is prepared by dissolv- 
ing barium carbonate in hydrochloric acid. It dissolves 
easily in water, but not as easily as the chlorides of 
strontium and calcium. The order of solubility, beginning 
with the most soluble, is, calcium, strontium, barium, — 
the same as in the case of the sulphates. 

Barium Hydroxide, Ba(OII) 2 , is formed by dissolving 
barium oxide in water, just as calcium hydroxide is formed 
by treating calcium oxide with water. In hot water it is 
much more easily soluble than calcium hydroxide, and it 
is also more easily soluble in cold water. As such a solu- 
tion acts in the same general way as lime-water, it is 
sometimes used in the laboratory for the purpose of 
detecting carbon dioxide, barium carbonate being insolu- 
ble. Like lime-water, it has an alkaline reaction. 

Barium Oxide, BaO, is made by heating the nitrate, as 
the carbonate is not easily decomposed by heat. The 
most interesting property of the oxide is its power to take 



512 COLLEGE CHEMISTRY. 

up oxygen when heated to dark red in the air or in 
oxygen, when it forms 

Barium Peroxide or Dioxide, Ba0 2 . — This peroxide is 
a white powder that looks like the simple oxide. When 
heated to a temperature a little higher than that required 
for its formation, it breaks down into barium oxide and 
oxygen. The formation of the peroxide by heating the 
oxide in the air, and the decomposition of the peroxide at 
a higher temperature, make it possible to extract oxygen 
from the air and to obtain it in the free state. This 
method of preparing oxygen on the large scale from the 
air was referred to under Oxygen. It is stated that the 
oxide improves with use. Specimens which have been in 
use for two years are said to be as efficient as at first. 
When a solution of hydrogen dioxide, H 2 2 , is added to a 
solution of barium hydroxide, a precipitate is formed which 
has the composition Ba0 2 -f- 8H 2 0. When filtered and 
put in a vacuum over sulphuric acid, it loses all its water 
and leaves behind pure dioxide. The dioxide is a con- 
venient starting-point in the preparation of hydrogen 
dioxide. It is only necessary to treat it with hydrochloric 
acid in order to make a solution of hydrogen dioxide. 
The solution made in this way, however, contains barium 
chloride. To make a solution containing nothing but the 
dioxide, pure barium peroxide is treated with dilute sul- 
phuric acid, when insoluble barium sulphate is formed and 
the hydrogen dioxide remains in solution : 

Ba0 2 + Ii 2 S0 4 = BaS0 4 + H 2 2 . 

It is interesting to compare the action of hydrochloric 
acid on barium peroxide and on the corresponding com- 
pound of manganese. As we have seen, with the latter 
the reaction takes place as represented in this equation : 

Mn0 2 + 4HC1 = MnCl, + 2H 2 + Cl 2 ; 

while with barium peroxide the reaction takes place thus : 

Ba0 2 + 2HC1 = BaCl 2 + H a O,. 



BARIUM COMPOUNDS. 513 

It is probable that in the case of manganese dioxide some 
intermediate reactions take place which are impossible in 
the other case. (See Manganese Dioxide.) 

Barium Sulphide, BaS, is made as calcium sulphide is, 
by reducing the sulphate by heating with charcoal. It is 
phosphorescent, like the calcium compound. 

Barium Sulphate, BaS0 4 . — This occurs in nature as 
barite, or heavy spar, and is precipitated when a soluble 
sulphate or sulphuric acid is added to a solution of a 
barium salt. It is insoluble in water; when freshly pre- 
cipitated, it is easily soluble in concentrated sulphuric 
acid. It is artificially prepared for use as a pigment and 
is known as permanent white. On account of its insolu- 
bility it is much used in chemical analysis for the purpose 
of detecting and estimating sulphuric acid. 

Barium Carbonate, Ba00 3 , occurs in nature as wither- 
ite, and is made pure by adding ammonium carbonate and 
a little ammonia to a solution of barium chloride. The 
carbonate usually found in the market is made by precipi- 
tating a solution of the crude sulphide with sodium car- 
bonate, or by heating together sodium carbonate and 
natural barium sulphate, or heavy spar. Made in either 
way it contains alkaline carbonate, from which it is im- 
possible to separate it by washing. The carbonate, like 
the other salts of barium, is poisonous. It has the power 
to unite, and form insoluble compounds, with metallic 
oxides of the formula M 2 0„ , as, for example, ferric oxide, 
Fe,0 3 , and is used in analytical operations for the purpose 
of separating iron from other metals, like manganese, 
which are not precipitated by it. 

Phosphates of Barium. — The phosphates of barium 
correspond in general to those of calcium. 

Reactions which are of Special Value in Analysis. 
— The sulphates of calcium and strontium are completely 
converted into the carbonates by contact with a solution 
of ammonium carbonate in ammonia. The sulphate of 
barium is not changed in this way. Consequently, if a 
mixture of the three sulphates is treated with ammonium 



5*4 COLLEGE CHEMISTRY. 

carbonate, those of calcium and strontium will be con- 
verted into carbonates, while that of barium will remain 
unchanged. By filtering, washing with water, and treat- 
ing with dilute nitric or hydrochloric acid, the carbonates 
will be dissolved, while the sulphate will not. If nitric 
acid is used, the solution may be evaporated to dryness, 
and treated with a mixture of alcohol and ether. Calcium 
nitrate will dissolve; strontium nitrate will not. 

Fluosilicic acid produces a precipitate of barium fluo- 
silicate, BaSiF 6 , in solutions of barium salts. This is 
insoluble in a mixture of alcohol and water, and difficultly 
soluble in water. The corresponding salts of calcium and 
strontium are soluble. 

Calcium sulphate in solution produces a precipitate in a 
solution of a strontium salt or of a barium salt, but not in 
one of a calcium salt. 

Strontium sulphate in solution precipitates barium sul- 
phate from a solution of a barium salt, but forms no pre- 
cipitate in a solution of a strontium salt. 

When boiled with a solution of one part of sodium 
carbonate and three parts of sodium sulphate, the sul- 
phates of strontium and calcium are completely convertc 1 
into carbonates, while the sulphate of barium remains 
unchanged. 

Barium chloride is insoluble in absolute alcohol ; calcium 
chloride is easily soluble; while strontium chloride dis- 
solves in warm absolute alcohol. 

Ammonium oxalate, (NH 4 ) 2 C 2 4 , produces precipitates 
of the oxalates in solutions of calcium, barium, and stron- 
tium. Only the calcium salt is insoluble in dilute acetic 
acid. 

Potassium bichromate, K 2 Cr 2 7 , precipitates barium 
chromate, BaCr0 4 . The corresponding salts of calcium 
and strontium are soluble in.water. Barium chromate is 
easily soluble in hydrochloric or nitric acid. 

All three elements of the group give colored flames 
which have characteristic spectra. Calcium compounds 
color the flame reddish yellow; strontium compounds give 



MAGNESIUM. 5 1 5 

an intense red; and barium compounds a yellowish green 
color. The spectra are more complicated than those of 
the elements of the potassium group, but each one contains 
highly characteristic lines which are easily recognized. 



MAGNESIUM SUB-GROUP. 
Glucinum, Gl (At. Wt. 9.1). 

Occurrence and Preparation. — The principal form in 
which the element glucinum occurs in nature is in the 
mineral beryl, which is a silicate of aluminium and 
glucinum of the formula Al 2 Gl 3 (Si0 3 ) 6 . Emerald has the 
same composition, but is colored green by the presence of 
a little chromic oxide. The element can be isolated by 
decomposing the chloride by heating it with potassium or 
sodium. 

Magnesium, Mg (At. Wt. 2±.36). 

Occurrence. — Magnesium occurs abundantly in nature, 
though by no means as abundantly as calcium. Among 
the widely distributed minerals that contain the element 
are magnesite, which is the carbonate, MgC0 3 ; dolomite, 
a double carbonate of magnesium and calcium: serpentine, 
talc, soapstone, meerschaum, hornblende, all of which 
contain magnesium silicates. Further, the metal is found 
in solution in many spring-waters in the form of the sul- 
phate, or, as it is called, Epsom salt. Kainite is a sul- 
phate and chloride of the composition expressed by the 
formula K 2 SO,. MgS0 4 . Mg01 2 + 6H 2 0: kieserite is mag- 
nesium sulphate, MgS0 4 -{-H 2 0; carnallite is a double 
chloride, KMgCl 3 + 6H 2 0. 

Magnesium compounds are contained in the soil in 
consequence of the decomposition of minerals containing 
it. It is to some extent taken up by the plants, and subse- 
quently into the animal body. It is found in the bones 
and in the blood in small quantities. 

Preparation. — The metal is made by the electrolysis of 
its chloride. In practice carnallite, MgKCl 3 -f- 6H 2 0, is 



516 COLLEGE CHEMISTRY. 

used. This is dehydrated and then melted in an iron 
crucible. A carbon rod brought into the molten mass 
serves as the anode, while the crucible itself acts as the 
cathode. The metal magnesium separates at the cathode. 

Properties. — It is a silver-white metal with a high lustre. 
In the air it changes only slowly, but it gradually becomes 
covered with a layer of the hydroxide. At ordinary tem- 
peratures magnesium does not decompose water: at 100° 
it decomposes it slowly. When heated above its melting- 
point in oxygen or in the air, it takes fire and burns with 
a bright flame, forming the white oxide. The light of the 
flame is very efficient- in producing certain chemical 
changes, such as those involved in photography, when a 
permanent impression is made by the light upon a sensitive 
plate. It has also the power to cause hydrogen and 
chlorine to combine just as the sunlight and the electric 
light do. 

Applications. — The principal use to which magnesium 
is put is for the purpose of producing a bright light, as 
for photographing in spaces to which the sunlight does 
not have access, and for signaling. It is also used to some 
extent as an ingredient of materials employed in making 
fireworks. 

Compounds of Magnesium. — The compounds of mag- 
nesium present a general resemblance to those of glucinum. 
As the element is much more abundant in nature, its 
compounds have been studied more extensively. Its acid 
properties are somewhat weaker, and its basic properties 
stronger, than those of glucinum. Its hydroxide does not 
form salts with the hydroxides of potassium and sodium. 
On the other hand, its chloride is decomposed when its 
water solution is evaporated to dryness. The hydroxide 
is very slightly soluble in water, and this solution has a 
slightly alkaline reaction. 

Magnesium Chloride, MgCl 2 .— This salt, as has been 
stated, occurs in nature. It is easily formed by dissolving 
magnesium oxide or carbonate in hydrochloric acid. On 
evaporating at as low a temperature as possible, there 



MAGNESIUM OXIDE— MAGNESIUM SULPHATE. 517 

finally crystallizes out of the very concentrated solution a 
salt of the composition MgCl 2 + 6ll 2 0, analogous to 
crystallized calcium chloride, CaCl 2 -J- 6H 2 0, and stron- 
tium chloride, SrCl, + 6H 2 0. When this crystallized salt 
is heated for the purpose of driving off the water, it is 
completely converted into the oxide 

MgCl 2 + H a O = MgO + 2HC1. 

The chloride is a white, crystalline mass that de- 
liquesces in the air. At a bright red heat, it is volatile 
and can be distilled in an atmosphere of hydrogen. 

Magnesium Oxide, MgO. — This compound is commonly 
called magnesia. A fine white variety, known as magnesia 
usta, is made by heating precipitated basic magnesium 
carbonate. It is a white, loose powder, which is difficultly 
soluble in water, forming with it the hydroxide, Mg(OH) 2 , 
which is also difficultly soluble. Magnesia is used, in 
medicine, as an application to wounds, and, mixed with a 
solution of ferric sulphate, as an antidote in cases of 
poisoning by arsenic. As magnesia is infusible, it is used 
to protect vessels which are subjected to a high tempera- 
ture. When mixed with water and allowed to lie in the 
air, it becomes very hard. Mixtures of magnesia with 
sand also have this property, and are used as hydraulic 
cements. It is used, further, in the manufacture of fire- 
bricks. 

Magnesium Sulphate, MgS0 4 . — The mineral kieserite, 
which occurs in Stassfurt, has the composition expressed 
by the formula MgS0 4 -f- H 2 0; or, more probably, this 
should be written (HO) 2 MgS0 3 , in which it appears as a 
derivative of the acid SO(OH) 4 . The salt, MgS0 4 + 7H 2 
(or H 2 MgS0 5 -j- 6H 2 0), also occurs in nature. It is this 
variety which is generally obtained when a solution of 
magnesium sulphate is evaporated to crystallization. It 
crystallizes in large rhombic prisms, or, if rapidly deposited 
from very concentrated solutions, in small, needle-shaped 
crystals. At ordinary temperatures, 100 parts of water 



518 COLLEGE CHEMISTRY. 

dissolve 125 parts of the salt. The water solution has a 
bitter, salty taste. When heated, it readily loses 6 mole- 
cules of water, but it requires a temperature of over 200° 
to drive oft' the last molecule. This has led to the belief 
that the salt with one molecule has the constitution above 
given, being a derivative of the acid SO(OH) 4 . 

Applications. — Magnesium sulphate finds extensive ap- 
plication. It is used in medicine as a purgative, and is 
known as Epsom salt, for the reason that it is contained 
in the water of Epsom springs; it is used further in the 
manufacture of sodium sulphate and potassium sulphate, 
and as a fertilizer in place of gypsum, it having been 
shown to be advantageous in some cases. Its chief use is 
for loading cotton fabrics. 

Magnesium Carbonate, MgC0 3 . — Magnesium shows a 
marked tendency to form basic salts with carbonic acid. 
When a neutral magnesium salt is treated with a soluble 
carbonate, a basic carbonate is precipitated, the composi- 
tion of which varies according to the conditions under 
which it is prepared. The salt obtained by adding an 
excess of sodium carbonate to a solution of magnesium 
sulphate has the composition, Mg 3 (OH) 2 (C0 3 ) 2 . The basic 
carbonate manufactured on the large scale is known as 
magnesia alba. It is this form of the carbonate that is 
used in the preparation of magnesia usta. Normal mag- 
nesium carbonate, MgC0 3 , occurs in nature as magnesite. 
It crystallizes in the same form as calcium carbonate, or 
is isomorphous with it. It is insoluble in water, but like 
calcium carbonate it dissolves in water containing carbon 
dioxide in solution. From this solution crystals having 
the composition MgC0 3 -j- 3H 2 and MgC0 3 + 5H 2 are 
deposited under the proper conditions. 

Phosphates. — The conduct of the phosphates of mag- 
nesium is very similar to that of the phosphates of calcium. 
All three are known; and of these only the primary salt is 
soluble in water. A salt much utilized in analysis is 
ammonium-magnesium phospliate, Mg(NH 4 )P0 4 . This is 
difficultly soluble in water, and may therefore be used 



BORATES, SILICATES, ETC. 519 

either for the purpose of detecting magnesium or phos- 
phoric acid. 

The corresponding salt of arsenic acid, Mg(NH 4 )As0 4 , 
is very similar to the phosphate, and on account of its 
insolubility it is also used in chemical analysis. 

Borates. — A borate of magnesium together with mag- 
nesium chloride occurs in nature, and is known as boracite. 
It has the composition expressed by the formula 2Mg 3 B 8 15 
-j- MgCl 2 . The borate, Mg 3 B 8 15 , is derived from the 
acid, H 6 B 8 15 , which is related to normal boric acid, as is 
shown by the equation 

8B(OH) 3 = H 6 B 9 15 + 9H 2 0. 

Silicates. — The simplest silicate of magnesium found in 
nature is olivine, which is represented by the formula 
Mg 2 Si0 4 . It is the neutral salt of normal silicic acid. 

Serpentine is derived from the acid, 0<n!LJ 3 , and has 

the composition Mg 3 Si 2 7 -{- 2H 2 0. 

Magnesium Silicide, Mg 2 Si, is made by heating together 
magnesium chloride, sodium fluosilicate, sodium chloride, 
and sodium. Under these circumstances the sodium sets 
magnesium free from the chloride, and silicon from the 
fluosilicate. Both unite to form magnesium silicide. 
When treated with hydrochloric acid it gives silicon 
hydride, SiH 4 , and hydrogen: 

Mg 2 Si + 4HC1 = 2MgCl 2 + SiH 4 . 

The liberation of hydrogen is due to the presence of an 
excess of magnesium. 

Reactions of Magnesium Salts which are of Special 
Value in Chemical Analysis. — Soluble hydroxides (KOH, 
NaOH, NH 4 OH) precipitate magnesium hydroxide. If 
ammonium chloride is present ammonia does not precipi- 
tate the hydroxide. ' 

Di- sodium phosphate with ammonia and ammonium 
chloride precipitates ammonium-magnesium phosphate 
from the solution of a magnesium salt. 



520 COLLEGE CHEMISTRY. 

Soclium and potassium carbonates precipitate basic 
magnesium carbonate. 

Erbium, E (At. Wt. 166). 

As regards the position of erbium in the periodic system, 
a final statement cannot as yet be made. It is even ques- 
tionable whether it is an element. 

EXPERIMENTS. 

Calcium Salts. 

Experiment 218. — Dissolve 10 to 20 grams of limestone or 
marble in common hydrochloric acid. Filter, and evaporate to 
dryness. Expose a few pieces of the residue to the air. 

Magnesium and its Salts. 

Experiment 219. — Make anhydrous magnesium chloride thus : 
Dissolve 180 grams magnesia usta in ordinary hydrochloric acid; 
shake the solution with an excess of magnesia to remove iron and 
aluminium ; filter ; add 400 grams ammonium chloride ; evapo- 
rate to dryness, keeping the mass constantly stirred. The double 
salt thus formed must be dried until a small specimen put in a 
test-tube is found not to give off water when heated. The dry 
salt is then ignited in a crucible placed in a furnace until ammo- 
nium chloride is no longer given off, when the molten mass, 
which is anhydrous magnesium chloride, is poured out on a stone 
and, after it is broken up, it is put in a dry bottle provided with 
a good stopper. 



CHAPTER XXVII. 

ELEMENTS OF FAMILY III, GROUP A: 

ALUMINIUM.— SCANDIUM.— YTTRIUM.— YTTERBIUM.— 

SAMARIUM.— HELIUM. 

General. — There is in some respects a resemblance 
between boron and the principal member of this group; 
but as boron acts almost exclusively as an acid-forming 
element, it was treated of in connection with the elements 
of Family V, Group B, or the nitrogen group. Attention 
was, however, called to the fact that the analogy between 
these elements and boron is but slight. The points of 
resemblance between boron and the members of Family 
III, Group A, will be indicated below. The principal 
member of this group is aluminium. The others are all 
rare, and some have been but imperfectly studied, owing 
to serious difficulties in the way of obtaining their com- 
pounds in pure condition* They are trivalent in their 
compounds, the general formulas being such as the follow- 
ing: 

Mfl\, M(OH) 3 , M(N0 3 ) 3 , M 2 (S0 4 ) 3 , M 2 (C0 3 ) 3 , MP0 4 , etc. 

Aluminium oxide is weakly basic, and somewhat acidic, 
though less so than boron oxide. Aluminium hydroxide 
has the power to neutralize most acids, and also to form 
salts with strong bases. Boron oxide, on the other hand, 
has scarcely any basic properties, though it does form a 
few extremely stable compounds, in which the boron takes 
the place of the hydrogen of acids. (See Boron Phos- 
phate, p. 342.) 

521 



522 COLLEGE CHEMISTRY. 

Aluminium, Al (At. Wt. 27.1). 

Occurrence. — Aluminium is an extremely important 
element in nature and in the arts. It occurs widely dis- 
tributed, and abundantly in many different forms of com- 
bination. Among them are feldspar, mica, cryolite, 
bauxite. Feldspar is a silicate of aluminium and potas- 
sium of the formula AlKSi 3 8 . Mica is a general name 
applied to a large number of minerals which are silicates 
of aluminium and some other metal, as potassium, lithium, 
magnesium, etc. The simplest form of mica is that repre- 
sented by the general formula KAlSi0 4 , according to 
which the mineral is a salt of orthosilicic acid, Si(OH) 4 . 
Cryolite is a double fluoride of aluminium and sodium, or 
the sodium salt of fluoaluminic acid, Na 3 AlF 6 . Bauxite is 
a hydroxide of aluminium in combination with a hydroxide 
of iron. Aluminium also occurs in the products of decom- 
position of minerals. One of the most important of these 
is clay, which is found in all conditions of purity — from 
the white kaolin to ordinary dark-colored clay. Kaolin 
is the aluminium salt of orthosilicic acid of the formula 
Al 4 (Si0 4 ) 3 -|- 4H 2 0. Aluminium silicate is found in all 
soils, but is not taken up by plants, and does not find 
entrance into the animal body. The name aluminium has 
its origin in the fact that the salt alum was known at an 
early date, and the metal was afterwards isolated from it. 

Preparation — The preparation of aluminium on the 
large scale presents a problem of the highest importance. 
The element has properties which adapt it to many uses to 
which iron is put, and for many purposes it has advantages 
over iron. Further, we are supplied by nature with un- 
limited quantities of the compounds of aluminium, which 
are distributed everywhere over the earth. While, how- 
ever, iron, lead, tin, copper, and other metals can be 
isolated from their natural compounds without serious 
difficulty, aluminium, which is more abundant than any 
of them, and in many respects more valuable than any of 
them, is locked in its compounds so firmly that it is only 



A LU MINIUM. 5 2 3 

by comparatively complicated and expensive methods that 
it can be isolated. 

The best method for the preparation of aluminium con- 
sists in the electrolysis of aluminium oxide, in the form 
of corundum. The oxide is contained in iron crucibles, 
and carbon rods are introduced into the molten mass. 
The iron crucible forms one electrode and the carbon rods 
the other. 

Properties. — The color of aluminium is like that of tin, 
and it has a high lustre. It is very strong, and yet 
malleable. It is lighter than most metals in common 
use, its specific gravity being 2.5 to 2.7 according to the 
condition, while that of iron is 7.8, that of silver 10.57, 
that of tin 7.3, and that of lead 11.37. It does not 
change in dry or in moist air; and in the compact form it 
does not act upon water even at elevated temperatures. 
It melts at about 700°, which is higher than the melting- 
point of zinc, and lower than that of silver. Hydrochloric 
acid dissolves it with ease, forming aluminium chloride. 
At the ordinary temperatures nitric and sulphuric acids 
do not act upon it; at higher temperatures, however, 
action takes place, and the corresponding salts are formed. 
It dissolves in solutions of the caustic alkalies, forming 
the so-called aluminates. It reduces many oxides when 
heated with them to a sufficiently high temperature; and 
is used in the preparation of boron and silicon. When a 
mixture of finely divided aluminium and ferric oxide is 
heated by means of a burning magnesium wire or by 
means of a cartridge made of magnesium powder and 
barium peroxide, the aluminium is oxidized at the expense 
of the ferric oxide, and an extremely high temperature is 
produced. This is the brightest white heat (about 3000°) 
which is sufficient to melt the iron and also the aluminic 
oxide which are formed. A method of welding has been 
devised by G-oldschmidt which is based upon this reaction. 

Applications. — The metal is used to a considerable 
extent in the preparation of ornaments, and of useful 
articles in which lightness is a matter of importance, as 



524 COLLEGE CHEMISTRY. 

in telescopes, opera-glasses, cooking-utensils, etc. An 
alloy with a small percentage of silver is used for the 
beams of chemical balances. Aluminium bronze, which 
is an alloy with copper, is also used quite extensively. It 
will be again referred to under Copper. 

Aluminium Chloride, A1C1 3 . — When aluminium hydrox- 
ide is dissolved in hydrochloric acid a solution of alu- 
minium chloride is formed, and from this solution a 
compound of the formula A101 3 -\- 6H 2 can be obtained 
in crystallized form. Like calcium and magnesium chlo- 
rides, this salt is deliquescent. When heated to drive off 
the water the salt conducts itself like magnesium chloride, 
but the decomposition into the oxide and hydrochloric 
acid takes place more easily than that of magnesium chlo- 
ride. The reaction is represented by the equation 

2A1C1 3 + 3H 2 = A1 2 3 + 6HC1. 

The dry chloride is prepared by the same method as that 
used in the preparation of silicon chloride and boron 
chloride, viz., by passing chlorine over a heated mixture 
of the oxide and carbon. The chloride, being volatile, 
sublimes, and is deposited in the cool part of the vessel, 
when pure, as a white, laminated crystalline mass. 
Generally, however, it is more or less colored in conse- 
quence of the presence of impurities. When exposed to 
the air it attracts moisture and gives off hydrochloric acid. 
It dissolves in water very easily, with a marked evolution 
of heat, but, from what was said above, it is evident that 
it cannot be obtained from this solution again by evapora- 
tion. It is volatile without change. The specific gravity 
of its vapor has been determined by different observers, 
and, unfortunately, with different results. According to 
Deville and Troost, it is such as to lead to the formula 
Al CI.. Nilson and Pettersson, however, have found it to 
correspond to that required by the formula A1C1 3 , their 
determinations having been made at a higher temperature 
than those of Deville and Troost. Still later determina- 
tions by Crafts again lead to the formula A1 2 C1 6 . In view 



ALUMINIUM HYDROXIDE— ALBUMINATES. 525 

of the conflicting evidence, the formula for aluminium 
chloride used here is the simpler one. By means of it and 
similar formulas for the other compounds of aluminium, 
the reactions of the element can be expressed somewhat 
more easily and probably just as truthfully as by means of 
the more complicated formula. 

Aluminium Hydroxide, Al(OH) 3 . — Normal aluminium 
hydroxide, Al(OH) 3 , occurs in nature as the mineral 
hydrargillite. It is precipitated from a solution of 
aluminium chloride by ammonia: 

A1C1 3 + 3NH 4 OH = Al(OH) 8 + 3NH 4 C1. 

Obtained by precipitation it forms a gelatinous mass, 
which is suggestive of starch -paste, and it is on this 
account extremely difficult to wash it completely free from 
the substances in the solution. It dries in the air, form- 
ing a gummy substance which has the composition 
Al(OH) 3 . When heated under proper conditions it loses 
water, and forms the compound A10 2 H : 

Al(OH) s = A10.0H + H 2 0. 

Ihis compound is found in nature as the mineral diaspore. 
If heated to a higher temperature it is converted into the 
oxide, A1 2 3 : 

2A1(0H) 3 = A1 2 3 + 3H 2 0. 

In the conduct of the chloride and of the hydroxide 
aluminium exhibits a certain resemblance to boron. The 
acidic character of the latter is, however, more strongly 
marked than that of the former. Boron chloride is more 
easily decomposed by water than aluminium chloride, 
and, as the decomposition takes place at the ordinary tem- 
perature, the product is the hydroxide instead of the 
oxide, as in the case of aluminium. The hydroxide, 
B(OH) 3 , readily loses water and forms metaboric acid, 
which in composition is analogous to diaspore; and at a 
higher temperature the oxide, B 2 3 , is formed. 

Aluminates.— When sodium or potassium hydroxide is 
added to a solution of an aluminium salt, aluminium 



526 COLLEGE CHEMISTRY. 

hydroxide is at first precipitated, but an excess of the 
reagent dissolves the precipitate. This action is due to 
the acidic character of aluminium hydroxide. It is prob- 
able that in solution the action with potassium and sodium 
hydroxides is of the same kind as represented in the equa- 
tions 

Al(OH) 3 + 3KOH = Al(OK) 3 + 3H 2 0, and 
Al(OH) 3 + 3NaOH = Al(ONa) 3 + 3H 2 0. 

On evaporating the solution of the potassium salt, how- 
ever, the product obtained has the formula AlO. OK, and 
is plainly the salt of the hydroxide AlO. OH, which may 
be called meta-aluminic acid, to suggest its analogy to 
metaboric acid, BO. OH. When aluminium hydroxide 
and sodium carbonate are melted together, the salt 
AlO.ONa is formed, as has been shown by determining 
the amount of carbon dioxide given off when a known 
weight of the hydroxide is employed. When, however, 
the solution of the hydroxide in caustic soda is evaporated, 
the salt Al(ONa) 3 is deposited. 

These salts are very unstable, though their solutions 
can be boiled without undergoing decomposition. Carbon 
dioxide decomposes them at once with precipitation of 
aluminium hydroxide. Similar salts are formed with 
calcium and barium. Among them may be mentioned 
those of the following formulas: Ca 3 (A10 3 ) 2 , Ca(A10 2 ) 2 , 
Ba 3 (A10 3 ) 2 , and Ba(A10 2 ) 2 . The calcium salts are insolu- 
ble in water, and some of them become hard in contact 
with water. They are therefore of importance in the 
manufacture of hydraulic cements. The barium salts are 
soluble in water. 

Many aluminates occur in nature, forming the impor- 
tant group of minerals known as the spinels. Of these, 
spinel itself is the magnesium salt of the hydroxide 

AlO. OH, and is represented by the formula .,^'^>Mg, 

or Mg( A10 2 ),. Chrysoberyl is the corresponding glucinum 
salt G1(A10 2 ) 2 ; and gahnite is the zinc salt Zn(A10 2 ) a . 



ALUMINATES. 5 2 7 

These salts are extremely stable, differing markedly in this 
respect from those above referred to, which are made in 
the laboratory. They are decomposed by heating them, 
in finely powdered condition, with primary or acid potas- 
sium sulphate, the action of which was described on 
p. 406. As will be seen farther on, there are other salts 
similar to the aluminates in structure which occur in 
nature. Among these chromic iron, or chromite, which 
is an iron salt of a hydroxide of chromium of the formula 
CrO.OH, may be mentioned here. The salt is to be 

regarded as made up according to the formula ,, ~' ^ > Fe, 

or Fe(Cr0 2 ) 2 . Further, magnetic oxide of iron or mag- 
netite, Fe 3 4 , is regarded as belonging to the same group, 

and its constitution represented thus: -™ ^'^>Fe, or 
Fe(Fe0 2 ) 2 ; and there is also a compound of magnesium, 
-p ( ^\>Mg. For the sake of emphasizing these anal- 
ogies, the formulas of the compounds above mentioned 
are here presented in tabular form : 

Potassium aluminate AIO.OK 

Sodium aluminate AlO.ONa 

Calcium aluminate AlO*0 > ^ a 

Barium aluminate AlO O^ ^ a 

Spmel A10.0 >Mg 

Chrysoberyll AlO 0> G1 

(ialmite . ,^ n > 

Chromite CrO <)> Fe 

Magnetite FeO.*0 >Fe 

Magnesio-ferrite p ^'^>Mg 



528 COLLEGE CHEMISTRY. 

Aluminium Oxide, A1 2 3 . — As has been stated, the 
oxide is formed by heating the hydroxide. It is found in 
nature in the form of ruby, sapphire, and corundum. 
The natural variety is extremely hard; and granular corun- 
dum, which is known as emery, is used for polishing. 
The red color of the ruby is caused by the presence of a 
trace of a chromium compound; while the blue color of 
the sapphire is probably due to the presence of a trace of 
a cobalt compound. Aluminium oxide is infusible in the 
hottest furnace fire, but it melts in the flame of the oxy- 
hydrogen blowpipe, and on cooling it becomes crystalline. 
In the electric furnace it melts and is vaporized. By 
mixing it with various easily fusible substances and heat- 
ing, it is obtained in the form of crystals, and by adding 
certain metallic oxides these crystals can be colored. In 
this way artificial rubies and sapphires have been prepared, 
which have all the properties of the natural ones. When 
the oxide is moistened with a few drops of a solution of 
cobaltous nitrate and then ignited, it turns blue. This 
fact is taken advantage of in chemical analysis for the 
purpose of detecting aluminium. When the oxide is 
made by gently igniting the hydroxide, it dissolves in 
strong acids. If, however, it is heated to a high tempera- 
ture, acids will not dissolve it. The natural varieties of 
the oxide, further, are not soluble in acids. By fusion 
with acid potassium sulphate insoluble aluminium oxide 
is converted into a soluble compound. 

Aluminium Sulphate, Al 2 (SOJ 3 . — This salt is made by 
dissolving the hydroxide of aluminium in dilute sulphuric 
acid, and evaporating to crystallization, when a salt of the 
composition Al 2 (SOJ 3 -f 18H 2 is deposited. When 
heated the salt loses its water of crystallization, and, if 
the temperature is raised to that of red heat, the anhydrous 
salt is decomposed with loss of sulphur trioxide and forma- 
tion of aluminium oxide. This decomposition is, however, 
not complete. The sulphate is manufactured on the large 
scale for various purposes, as, for example, for a mordant, 
for sizing paper, etc. 



ALUMS. 529 

Alums. — When a solution of aluminium sulphate is 
brought together with a solution of potassium sulphate in 
the proportion of their molecular weights, a salt crystal 
lizes out which has the composition represented by the 
formula 

KA1(S0 4 ) 2 + 12H 2 or K 2 S0 4 + A1 2 (S0 4 ) 3 + 24H 2 0. 

The most rational view held regarding this compound is 

SO^ 0K 

b0 2<(X 

that it has the constitution QA 0— >A1, with perhaps 

some of the so-called water of crystallization present in 
the form of hydroxy!. This salt, which has long been 
known under the name of alum, is the type of a class of 
similar compounds, all of which are called alums. These 
may be regarded as derived from the ordinary form by 
substituting sodium, ammonium, or any other member of 
the sodium group, besides some other metals, for the 
potassium. Thus a series of alums is obtained, of which 
the following are examples: 

NaAl(S0 4 ) 2 + 12H 2 0; 
LiAl(S0 4 ) 2 + 12H 2 0; 
(NH 4 )A1(S0 4 ) 2 + 12H,0; 
CsAl(S0 4 ) 2 + 12H 2 0. 

Again, alums are derived from the ordinary form by sub- 
stituting for the aluminium some other elements that 
resemble aluminium, as, for example, iron, chromium, 
and manganese. Such alums are those represented by the 
following formulas : 

KFe(S0 4 ) 2 + 12H 2 0; 
KCr(S0 4 ) 2 +12H 2 0; 
KMn(S0 4 ) 2 + 12H 2 0. 

All the alums have certain properties in common. They 
are all soluble in water, and all crystallize in the same 
forms, which are regular octahedrons combined with 



53° COLLEGE CHEMISTRY. 

cubes. If a crystal of one alum is suspended in the solu- 
tion of any other one it will continue to grow. They are 
all strictly isomorphous. The principal alums containing 
aluminium are those of potassium and ammonium, both 
of which are manufactured on the large scale. 

Potassium Alum, Potassium - Aluminium Sulphate, 
KA1(S0 4 ) 2 -f 12H a C— Ordinary alum is found in nature 
in some volcanic regions. The mineral alunite, which is 
a basic salt of the formula K(A10) 3 (S0 4 ) 2 + 3H 2 0, or per- 
haps K[A1(0H) 2 ] 3 (S0 4 ) 2 , occurs in larger quantities. 
When this salt is heated and treated with water, ordinary 
alum dissolves, and is easily obtained from the solution. 
Another source of alum is alum shale. This occurs in 
large quantities in nature, and consists of coal, clay, and 
iron pyrites. When it is heated in contact with the air 
the coal burns, as do also the sulphur and pyrites, and 
sulphuric acid is formed. When allowed to lie for a time 
in contact with the air the iron pyrites is converted into 
sulphate and sulphuric acid. The latter acts upon the 
clay or aluminium silicate, forming aluminium sulphate, 
from which alum can easily be made. It is easier to treat 
the shale and similar substances with sulphuric acid 
directly, and this method is now generally employed. 

Alum dissolves readily in hot water, 357.5 parts of the 
crystallized salt dissolving in 100 parts of water at 100°. 
At 0° only 3.9 parts dissolve, and at the ordinary tempera- 
ture about 12 parts. It crystallizes beautifully in regular 
octahedrons, occasionally with cube faces developed on 
them. Under some circumstances it crystallizes in cubes. 
When heated, alum melts in its water of crystallization, 
and if heated to a sufficiently high temperature the water 
passes off, leaving burnt alum. 

Applications. — Alum is used very extensively in the 
preparation of pigments, as a mordant, in the sizing of 
paper, for clarifying water, etc. 

Ammonium Alum, Ammonium-Aluminium Sulphate, 
(NH 4 ) A1(S0 4 ) 2 -f- 12H 2 0, is in most respects much like the 
potassium compound, and can be used in place of it for 



ALUMINIUM SlUC/tTES. 53 1 

almost all purposes for which alum is used. It is some 
what more easily soluble than ordinary alum. As it is 
cheaper than the latter it is largely manufactured in place 
of it. 

Sodium Alum is much more easily soluble in water than 
either potassium or ammonium alum, and this makes it 
difficult to prepare it in pure condition. It is therefore 
not manufactured, although sodium compounds are 
cheaper than those of either potassium or ammonium. 

Aluminium Silicates. — It has been stated that alumin- 
ium silicate enters into the composition of a number of 
important minerals. It occurs in enormous quantities in 
nature. The most important of the minerals containing 
it are the feldspars, of which ordinary feldspar, KAlSi 3 8 , 
and albite, or sodium feldspar, NaAlSi 3 8 , are the most 
abundant. These again enter into the composition of 
granite together with quartz and mica, and mica is itself 
a double silicate of aluminium. As remarked under Silicic 
Acid (which see), the natural silicates are for the most 
part salts of polysilicic acids which are derived from ortho- 
silicic acid by loss of water from two or more molecules. 
Up to the present but little more has been done with the 
many natural silicates than to determine their percentage 
composition. It appears probable from what has already 
been learned regarding their constitution that investiga- 
tions in this direction will before long yield interesting 
results. As yet, however, the methods for such investiga- 
tions are quite unsatisfactory, owing largely to the fact 
that the compounds are so extremely stable that but few 
reagents decompose them, and if they are decomposed at 
all, the products are such that no conclusion can be drawn 
regarding the constitution. A careful study of the rela- 
tions in which minerals occur in nature will undoubtedly 
be of assistance, as this will throw some light upon the 
conditions under which they were formed. 

One of the most common decompositions of minerals* 
constantly taking place, and that has taken place to an 
enormous extent, is that of feldspar. Under the influence 



S3 2 COLLEGE CHEMISTRY. 

of moisture and the carbon dioxide of the air, this sub- 
stance slowly decomposes, the products being mainly 
potassium or sodium silicate and aluminium silicate. The 
salt of the alkali metal, principally potassium, being solu- 
ble, is carried away, and finds its way into the soil. The 
silicate of aluminium is not soluble, bat it easily forms 
emulsions with water, and is therefore carried down the 
sides of the hills and mountainsupon which it is formed 
into the valleys, and much of it finds its way into streams. 
Sometimes this carrying away is prevented, and then large 
beds of comparatively pure clay, known as kaolin, are 
formed. The clay found in the valleys is always more or 
less impure and colored. 

Kaolin. — This is the purest form of aluminium silicate 
found in nature. It always contains water. Its composi- 
tion varies, some specimens on analysis giving results 
which lead to the formula Al 4 (SiOJ 8 + 4H 2 0, according 
to which the substance is the salt of normal silicic acid, 
Si(OH) 4 . Other specimens have the composition HAlSi0 4 
OH 

-(- H 2 0, or Si-j /-)\ ai _i_ TT O* When heated alone kaolin 

J)/ 
does not melt; but if feldspar is added to it, the whole 
melts, and forms a translucent mass known as porcelain. 
Other substances besides feldspar may be used for this 
purpose. 

Clay. — Clay, as has been stated, is a name given to cne 
impure varieties of aluminium "silicate which have been 
carried down from the place of formation. Among the 
substances besides aluminium silicate found in clays are 
calcium carbonate, magnesium carbonate, sand, and 
hydroxides of iron. The color is largely determined by 
the amount of the hydroxides of iron present. The better 
varieties are used in the manufacture of the so-called 
"stone -ware," gas-retorts, and fire-bricks. The colored 
varieties are used for making ordinary earthenware and 
bricks. Marl is clay mixed with considerable quantities 
of calcium carbonate. 



PORCELAIN. 533 

Ultramarine. — The substance occurring in nature and 
known as lapis lazuli consists of a silicate of sodium and 
aluminium together with a sulphur compound, probably a 
polysulphide of sodium. The coloring matter, known as 
ultramarine, obtained by powdering it was formerly very 
expensive, but it is now made artificially by the ton, and 
the color of the artificially prepared substance is even more 
beautiful than that of the natural. 

Porcelain. — It was stated above that when kaolin is 
heated alone it does not melt, but that when feldspar is 
added to it, or when that found in nature contains feld- 
spar, as is frequently the case, it either fuses, forming a 
compact mass, or it melts and forms a translucent mass. 
Further, when kaolin or any other variety of clay is mixed 
with water, a plastic substance results, that can be kneaded 
and worked into any desired form. These facts form the 
basis of the manufacture of earthenware, porcelain, etc. 
The ease with which the mass melts depends upon the 
quantity of feldspar or other flux added to it. If but little 
is added it melts with difficulty; if much is added it melts 
easily. 

In the manufacture of the finest kinds of porcelain 
kaolin is used. This is generally mixed with a little feld- 
spar or chalk, gypsum or some other flux, and sand is also 
added, All these substances must be very finely ground. 
The mixture is then worked into the desired forms, and 
carefully dried. After the objects are dried they are next 
burned, first at a red heat at which the mass becomes 
solid, afterwards at a white heat for the purpose of form- 
ing a glaze upon the surface. The product after the first 
burning is that which is familiar as porous earthenware; 
that formed in the second burning is the porcelain with 
glaze as it is commonly used. 

In order to form the glaze upon the porcelain the porous 
earthenware first formed is drawn through a vessel con- 
taining proper materials in finely powdered condition and 
suspended in water. The materials used are generally the 
same as those used for the porcelain itself, but they are 



534 COLLEGE CHEMISTRY. 

mixed in different proportions, with less kaolin, and more 
sand and feldspar, so as to make them more easily fusible. 
After this treatment the objects are again heated to a high 
temperature. 

Earthenware. — The ordinary varieties of earthenware are 
made from varieties of clay which are much less pure than 
kaolin. Ordinary colored clay is used. The objects are 
formed, and then subjected in general to the same kind of 
treatment as porcelain. They are glazed in different ways. 
One method consists in bringing the glazing material on 
the earthenware before it is burned; another method con- 
sists in putting the objects in the furnace without a glaze, 
and towards the end of the firing process sodium chloride 
is thrown into the furnace, and is thus brought in contact 
with the ware in the form of vapor. A chemical change 
takes place, resulting in the formation of a silicate of 
aluminium and sodium upon the surface. This melts, and 
forms a glaze. 

Bricks are the most common variety of unglazed 
earthenware. Owing to the presence of other substances 
besides aluminium silicate, as, for example, calcium car- 
bonate, the material is comparatively easily fusible. The 
color of bricks is largely due to the presence of oxides of 
iron. 

Alloys. — Aluminium forms alloys with magnesium. 
Those containing 10 to 25 per cent magnesium come into 
the market under the name of magnalium. This has a 
specific gravity 2 to 2.5, and is therefore lighter than 
aluminium, but it is much harder. Magnalinm containing 
a considerably higher percentage of aluminium is used in. 
making mirrors, for which it is well adapted on account of 
its hardness and its power of taking a high polish. 
Aluminium amalgam is an extremely active substance. It 
decomposes water energetically. 

Reactions of Aluminium Salts which are of Special 
Value in Chemical Analysis. — Potassium and sodium 
hydroxides precipitate aluminium hydroxide, which is 
soluble in an excess of the reagents. 



SCANDIUM. 535 

Ammonia precipitates the hydroxide, which is only 
slightly soluble in tin excess of the reagent. 

Hydrogen sulphide and carbon dioxide precipitate 
aluminium hydroxide from a solution of an aluminate; 
that is, from a solution of aluminium hydroxide in a 
caustic alkali. 

Ammonium sulphide and other soluble sulphides precipi- 
tate the hydroxide. This is due to the instability of the 
sulphide of aluminium, or, going farther back, to the 
weak basic character of the hydroxide. The reaction of 
ammonium sulphide with aluminium sulphate takes place 
as represented in the following equation : 

Al 2 (SO«)3 + 3(NH 4 ) 2 S + 6H 2 = 3(NH 4 ) 2 S0 4 + 3H 2 S + 2A1(0H) 3 . 

Soluble carbonates precipitate aluminium hydroxide for 
the same reason that the soluble sulphides do. The 
reaction between aluminium sulphate and sodium car- 
bonate takes place thus : 

A1 2 (S0 4 ) 3 + 3Na 2 C0 3 + 3H 2 = 3Na 2 S0 4 + 3C0 2 + 2A1(0H) 3 . 

Other Members of Family III, Group A. 

Scandium, Sc (At. TVt. 44.1). — This element was dis- 
covered in 1880 in the minerals euxenite and gadolinite. 
Its compounds are similar to those of aluminium. It 
forms an oxide of the formula Sc 2 3 ; a sulphate, Sc 2 (SOJ 3 ; 
a double sulphate, KSc(S0 4 ) 2 ; etc. It is of special interest 
for the reason that its properties were foretold by Men- 
deleeff several years before it was discovered. The 
prophecy was based upon the position of the element in 
the periodic system. When the relations between the 
atomic weights and properties of the elements were first 
described in a comprehensive way by Lothar Meyer and 
Mendeleeff, the latter described the properties of an ele- 
ment then unknown, which he called ekaboron. This 
should have the atomic weight about 44, should form an 
oxide of the formula M 2 3 , etc. It has been shown that 
the properties of scandium agree very closely with those 
foretold. 



53 6 COLLEGE CHEMISTRY. 

Yttrium, Y (At. Wt. 89), like scandium, is found in 
gadolinite, euxenite, and some other rare minerals. The 
element itself has not been isolated. 

Ytterbium, Yb (At. Wt. 173).— This rare element, like 
scandium and yttrium, is found in gadolinite and euxenite 
— most abundantly in the latter. Its compounds in 
general resemble those of yttrium. 

Samarium, Sin, has been obtained from samarskite, a 
North Carolina mineral. It has also been found in small 
quantities in some other minerals, as cerite, gadolinite, 
and orthite. In general it resembles aluminium in its 
chemical conduct. 

Helium. — When cleveite and some other rare minerals 
are treated with acids a gas is given off. This was at first 
held to be nitrogen, but a careful spectroscopic examina- 
tion by Kamsay has shown that it contains a hitherto 
unknown gas. The spectrum contains lines which are 
identical with lines that have been observed in the solar 
spectrum. These have been ascribed to an unknown 
element, helium. This has now been isolated and studied. 
It proves to be a very light gas, as incapable of forming 
chemical compounds as argon. Practically all the evidence 
goes to show that it is an element, but as yet little is 
known in regard to its relations to other elements. 

The chemistry of lanthanum is so intimately connected 
with that of cerium and didymium, that, although these 
three elements appear to belong to different families, they 
will be briefly treated of together. 

EXPERIMENT. 
Aluminium Chloride. 
Experiment 220. — Aluminium chloride is made thus : Mix 
aluminium oxide with starch-paste ; form the mass into small 
balls the size of ordinary marbles ; ignite these in a crucible in a 
furnace ; put them in a porcelain tube, and then pass dry chlo- 
rine over them, at the same time heating the tube to redness. 
The chloride will sublime and be deposited in the front end of the 
tube' or in a receiver if the heat is sufficient. It can be purified 
by subliming it over heated iron or aluminium. 



CHAPTER XXVIII. 

ELEMENTS OF FAMILY I, GROUP B: 
COPPER.— SILVER.— GOLD. • 

General. — The facts which strike one most forcibly on 
comparing the elements of this group with those of 
Group A of the same family are, that they are much less 
active chemically, and that they furnish a greater variety 
of compounds. Sodium and potassium and the other 
members of Group A display the greatest activity, as we 
have seen. The basic character is most strongly developed 
in them. Further, in nearly all their compounds they act 
with the same valence. They are univalent in all their 
salts. Copper, silver, and gold, however, are not chemic- 
ally active elements, and the activity grows less with 
increasing atomic weight. Copper and gold form two 
series of compounds each, and silver also forms a few com- 
pounds, in which it appears with a valence greater than 
one. In the two series of salts formed by copper tho 
element appears to be univalent and bivalent, as in the 
chlorides CuCl and CuCl 2 . Gold, however, is univalent 
and trivalent, while silver is almost exclusively univalent. 
It must be said that the resemblance between gold and 
the other members of Group B, Family I, is apparently 
not as marked as that between mercury and copper and 
silver. 

Copper, Cu (At. Wt. 63.6). 

General. — The compounds of copper which are most 
commonly met with are those in which it acts as a bivalent 
element. Its principal comoounds are copper oxide, CuO; 

537 



53^ COLLEGE CHEMISTRY. 

the sulphate, CuS0 4 ; and the sulphide, CuS. In all these 
the copper is bivalent. But besides these there are such 
compounds as CuOl and Cu 2 0, in which the element 
appears to be univalent. There are, then, two series of 
salts, of which the following will serve as examples : 



CuCl 


CuCl 2 


CuBr 


CuBr 2 


Cu 2 


CuO 



Those compounds which are of the first order, corre- 
sponding to the chloride CuCl, are called cuprous com- 
pounds. Thus, CuCl is cuprous chloride; Cu 2 0, cuprous 
oxide, etc. On the other hand, compounds of the second 
order are called cupric compounds. Thus, CuCl 2 is cupric 
chloride; CuO, cupric oxide; CuS0 4 , cupric sulphate, etc. 

Forms in which Copper occurs in Nature. — Copper is a 
widely distributed element, and it occurs also in large 
quantities. It occurs in the uncombined condition, or as 
native copper, in large quantity in the United States in the 
neighborhood of Lake Superior, in China, Japan, Siberia, 
and Sweden. The most valuable ores of copper are the 
oxides, ruby copper or cuprous oxide, Cu 2 0, and cupric 
oxide, CuO; the carbonates, as malachite, Cu 2 (OH) 2 C0 3 ; 
the sulphides, as chalcocite, Cu 2 S; copper pyrites, 
Cu 2 S.Fe 2 S 3 ; and others. 

Metallurgy of Copper. — The metallurgy of copper is 
comparatively complicated, owing to the difficulty of con- 
verting the ores of copper into the oxide. In most of the 
ores used sulphur and iron are contained, as well as 
smaller quantities of other elements, as arsenic, antimony, 
lead, etc. The ores are first roasted with the object of 
converting the sulphides partly into oxides. Under these 
circumstances the sulphides of iron are more easily con- 
verted into the oxides than the sulphides of copper. By 
adding a material rich in silicic acid, and melting the 
roasted ore in a blast furnace with charcoal, the oxide of 
iron is partly reduced, and converted into silicate, which 
runs off with the slag. In this way a product is obtained 



COPPER. 539 

which is richer in copper than the roasted ore. This, 
which is called the matte, contains copper sulphide and 
iron sulphide. The matte is again roasted and melted in 
the same way as the ore, and a further quantity of iron is 
removed, while some of the copper is reduced. A reaction 
which plays an important part in these processes is that 
which takes place between cuprous oxide and cuprous sul- 
phide, forming metallic copper and sulphur dioxide: 



2Cu 2 + Cu 2 S = 6Cu + SO 



Sometimes it is necessary to repeat the roasting and melt- 
ing with charcoal and sand a number of times, the matte 
becoming richer in copper at each successive stage. 

Properties. — Copper is a hard metal, of a reddish color 
and metallic lustre. It does not change in dry air, but in 
moist air it gradually becomes covered with a green layer of 
a basic carbonate. It melts at a somewhat lower tempera- 
ture than gold, and at a somewhat higher temperature than 
silver. It is very malleable and tenacious. It decomposes 
water only at bright-red heat. When heated in the air to 
a comparatively high temperature it becomes covered with 
a layer of cupric oxide; at a lower temperature cuprous 
oxide is formed. Nitric acid dissolves it, copper ni- 
trate, Cu(]Sr03) 2 , being formed, and the oxides of nitrogen 
being evolved (see p. 286); dilute sulphuric acid does 
not act upon it unless the air has access to it; concen- 
trated sulphuric acid when heated with it forms cupric 
sulphate, CuS0 4 , and sulphur dioxide. Dilute acids in 
general do not act upon it unless the air has access to 
it. This fact is of importance in connection with the use 
of copper vessels in culinary operations. Substances con- 
taining vegetable acids can be boiled in bright copper 
vessels with impunity, for the water-vapor prevents the 
access of air, but, on cooling, the air is admitted, and then 
action takes place, causing solution of some of the copper, 
which is objectionable. Ammonia in contact with copper 
absorbs oxygen, and the copper dissolves in consequence 
of the formation of a compound of cupric oxide and am- 



54° COLLEGE CHEMISTRY. 

mouia. This fact is sometimes taken advantage of for the 
preparation of nitrogen, as has already been stated (see p. 
251). 

Applications. — As is well known, copper is used very 
extensively for a variety of purposes, among which the 
following may be mentioned : for electrical apparatus, 
coins, copper vessels, roofs, for covering the bottoms of 
ships, etc. It is also used for copper-plating; and in the 
preparation of a number of valuable alloys, such as brass, 
bronze, gun-metal, bell-metal, etc. 

Alloys. — Brass is a mixture or compound of about 
one part of zinc and two parts of copper; these proportions 
may, however, be varied between quite wide limits. There 
is a variety of brass containing equal parts of zinc and 
copper, and another containing one part of zinc and five 
parts of copper. Pinchbeck is made by combining two 
parts of copper and one of brass. 

Bronze consists of copper, zinc, and tin. The propor- 
tion of copper varies from 65 to 84 per cent; that of zinc 
from 11 to 31.5 per cent; and that of tin from 2.5 to 4 
per cent. When exposed to the air bronze becomes 
covered with a green coating of basic copper carbonate, 
which protects it from further action. This coating is 
now generally produced artificially by a variety of methods, 
as by washing the surface with a solution of salts and 
acids. 

Gun-metal consists generally of copper and tin in the 
proportion of 11 parts of tin and 100 parts of copper. 

Bell-metal contains a larger proportion (from 20 to 25 
per cent) of tin than gun-metal. 

Alloys with Aluminium, containing aluminium and 
copper in widely different proportions are made. That 
with 3 per cent of copper has a whiter color than alumin- 
ium, the color being more like that of silver. On the 
other hand, an alloy of copper with 5 to 10 per cent of 
aluminium has a color resembling that of gold. This, 
which is known as aluminium bronze, is very hard and 
elastic, and is not easily acted upon by chemical reagents. 



COMPOUNDS OF COPPER. 54* 

It is now used to a considerable extent in the manufacture 
of ornamental and useful articles. 

German silver is an alloy consisting of copper, zinc, and 
nickel. The proportion of copper varies from 40 to GO 
per cent; that of zinc from 19 to 44 per cent; and that of 
nickel from to 22 per cent. 

Cuprous Chloride, CuOl, is formed by heating cupric 
chloride, CuCl 2 ; and by boiling an excess of copper filings 
with concentrated hydrochloric acid with the addition of 
a little nitric acid, filtering through asbestos, and pouring 
into water. It is a white crystalline compound, and is 
difficultly soluble in water. When exposed to the air it 
turns green in consequence of the formation of a basic 
chloride. It is volatile at a high temperature, and a 
determination of the specific gravity of the vapor gave a 
result corresponding to the formula Cu 2 Cl 2 . It has 
markedly the power to combine with chlorine, and there- 
fore acts as a reducing agent. 

Cupric Chloride, CuCl 2 . — This compound is formed by 
treating copper or cuprous chloride with chlorine. It is 
also easily made by dissolving cupric hydroxide, or car- 
bonate, in hydrochloric acid. From its solution in water 
the chloride crystallizes with two molecules of water, 
CuCl 2 + 2H 2 0. The crystals when heated lose their water 
without suffering further decomposition, except at high 
heat, when a part of the chlorine is given off, and cuprous 
chloride is formed. Cupric chloride combines with am- 
monia gas, forming a compound of the formula CuCl 2 . 6NH 3 
which is soluble in water, with a dark blue color. When 
heated it loses four molecules of ammonia, and the com- 
pound CuCl 2 .2KB 3 is left behind. 

Cuprous Hydroxide, Cu(OII). — The simple compound 
of the formula here given is not known, but a derivative 
of this, of the formula Cu 8 3 (OH) 2 or 40u 2 O.H 2 0, is easily 
made by adding sodium hydroxide to a solution of a 
cuprous salt. It passes readily over into cuprous oxide 
when gently heated. 

Cuprous Oxide, Cu 2 0, occurs in nature, and is known 



542 COLLEGE CHEMISTRY. 

as ruby copper or cuprite. It is easily prepared by treating 
a solution of glucose, or starch sugar, with copper sulphate 
and potassium hydroxide. By boiling, the copper is 
thrown down in the form of cuprous oxide. At first this 
is yellow, and it is supposed by some that the yellow com- 
pound is the hydroxide, but satisfactory evidence of this 
has not been furnished. The yellow precipitate is soon 
converted into the red oxide. Cuprous oxide is not 
changed when allowed to lie in contact with the air. It 
dissolves in nitric and sulphuric acids, forming cupric 
salts; and if the acids are dilute, copper is deposited. 
This will be clear from a consideration of the following 
equation : 

Ou 2 + H 2 S0 4 = CuS0 4 + H 2 + Cu. 

Cupric Hydroxide, Cu(OH) 2 , like the hydroxides of 
most base-forming elements, is thrown down by the addi- 
tion of a soluble hydroxide to a cupric salt. It is a 
voluminous, blue precipitate. When allowed to stand in 
a solution, or when the solution is boiled, the hydroxide 
loses a part of its hydroxyl, and is converted into a black 
compound of the formula Ou(OH) 2 -j- 2CuO, and this when 
dried and heated is converted into the oxide CuO. 

Cupric Oxide, CuO. — Cupric oxide is found in nature 
in the neighborhood of Lake Superior in the United States, 
and is formed by heating copper to redness in contact with 
the air, or by heating the nitrate. It loses its oxygen very 
readily when -treated with reducing agents, such as 
hydrogen and carbon. It is used extensively in quantita- 
tive analysis for the purpose of estimating the composition 
of organic compounds, or such as contain carbon and 
hydrogen. Its use is based upon the fact that when 
organic compounds are heated with the oxide they are 
oxidized, the carbon being converted into carbon dioxide 
and the hydrogen into water. By passing the products of 
the oxidation through calcium chloride, and a solution of 
potassium hydroxide, the water is retained in the first, 



CUPRIC SULPHATE. 543 

and the carbon dioxide in the second, and the weight of 
each formed can easily be determined. 

Cupric Sulphate, CuS0 4 . — This salt is manufactured on 
the large scale, and in the crystallized form, containing five 
molecules of water, CuS0 4 -f 5H 2 0, is commonly called 
'•blue vitriol " or "blue stone." It is found, to some 
extent in nature, being formed by the action of the oxygen 
of the air on the sulphide. It is most conveniently made 
by dissolving metallic copper in concentrated sulphuric 
acid, or by treating cupric sulphide with sulphuric acid. 
The action of sulphuric acid on the metal consists essen- 
tially in the formation of cupric sulphate, sulphur dioxide, 
and water, as expressed in the equation 

Cu + 2H 2 S0 4 = CuS0 4 + S0 2 + 2H 2 0. 

The question whether the copper reduces the sulphuric 
acid directly, or the hydrogen given off from the acid 
effects the reduction, is an open one. But there are other 
products formed besides those mentioned. At first a 
brown substance of the composition Cu 2 S is deposited. 
As the action proceeds oxysulphides are formed, the final 
product of a series of changes being Cu 2 OS, or CuO.CuS, 
which is black and insoluble in water. Under some con- 
ditions a considerable proportion of the copper is trans- 
formed into the oxysulphides by sulphuric acid. Cupric 
sulphate is obtained in large blue crystals of the triclinic 
system, which have the composition CuS0 4 -|- 5H 2 0. 
When heated to 100°, four molecules of water are given 
off, and the last is not given off until the temperature 
200° is reached. This makes it appear probable that the 
salt has the constitution represented by the formula 
CuS0 3 (OH) 2 , corresponding in this respect to magnesium 
sulphate (which see). When heated higher, it loses all its 
hydroxy!, and the salt, CuS0 4 , is left in the form of a 
white powder, which has the power to take up water from 
the air, becoming blue again. It dissolves in three parts 
oi cold water and one-half part boiling water. Copper 



544 COLLEGE CHEMISTRY. 

sulphate, containing seven molecules of water, CuS0 4 -{- 
7H 2 0, is obtained when mixed with solutions of the sul- 
phates of iron, zinc, or magnesium, all of which crystallize 
with seven molecules of water. In this form cupric sul- 
phate is isomorphous with the other sulphates. These 
salts have in general received the name of vitriols, and the 
old names " green vitriol," "white vitriol," and "blue 
vitriol " are still used to some extent, though rarely by 
chemists. Among the similar salts included under the 
same general head are the following : 

Zinc sulphate (white vitriol) ZnS0 4 -f 7H 2 

Magnesium sulphate MgS0 4 -f 7H 2 

Glucinum sulphate G1SO, + 7H 2 

Ferrous sulphate (green vitriol) FeS0 4 -f 7H 2 

Nickel sulphate NiS0 4 + 7H 2 

Cobalt sulphate CoS0 4 + 7H s O 

Copper sulphate (blue vitriol) CuS0 4 + 7H 2 

(CuS0 4 + 5H 2 0) 

Applications. — Cupric sulphate is used extensively in 
the preparation of. blue and green pigments, in copper- 
plating by electrolysis, in galvanic batteries, for the pur- 
pose of preserving wood, and as a remedy against phylloxera 
(see p. 382), etc. 

Cupric Nitrate, Cu(N0 3 ) 2 , is easily formed by dissolving 
copper in dilute nitric acid. It is easily soluble in water, 
and is deposited in crystallized form, the crystals contain- 
ing three or six molecules of water according to the tem- 
perature, the salt with six molecules being formed at the 
lower temperature. Like other copper salts, it has a blue 
color. 

Cupric Arsenite, CuHAs0 3 , is formed as a greenish- 
yellow precipitate when an ammoniacal solution of arseni- 
ous acid is added to a solution of cupric sulphate. It is 
known as Scheele's green. A compound of cupric arsenite 
and cupric acetate, which is made by treating a basic 
acetate of copper with arseriious acid, is known as Paris 
green or Schiveinfurt green. 



CUPRIC CARBONATES, ETC. 545 

Cupric Carbonates. — When a soluble carbonate is 
added to a solution of cnpric sulphate a voluminous 
greenish precipitate is formed, which has the composition 

Cu< OH 
Cu (OH) r CuC0 3 , or q > CO. This is plainly a basic 

c«< 0H 

carbonate. The mineral malachite, which has a beautiful 
green color, has the same composition as the precipitate 
just mentioned. 

Cuprous Sulphide, Cu a S.— This compound occurs in 
nature, and is known as chalcocite. It is, further, a con- 
stituent of copper pyrites, which is a compound of cuprous 
and ferric sulphides, Cn 3 S.Fe 2 S 3 or CuFeS 2 . It can be 
made by heating copper and sulphur together in the right 
proportions. It has a grayish-black color; does not give 
up its sulphur, even when heated in hydrogen; and is the 
more stable of the two sulphides of copper. 

Cupric Sulphide, CuS. — This is formed as a black pre- 
cipitate when hydrogen sulphide is passed into a solution 
of a cupric salt. In water alone cupric sulphide is some- 
what soluble. Hence in washing out a precipitate of 
copper sulphide with water a little of it will pass through 
in solution. It also easily undergoes oxidation, and, as it 
forms the sulphate, some is dissolved in this way unless 
proper precautions are taken. It is slightly soluble in 
ammonium sulphide, but insoluble in sodium sulphide. 
The above facts are of importance in the analysis of com- 
pounds containing copper, as will readily be seen. When 
heated, cupric sulphide loses half its sulphur, and is con- 
verted into cuprous sulphide. 

Copper-plating. — The process of copper-plating consists 
in brief in depositing upon an object a layer of copper 
by putting it in a bath containing some copper salt, and 
connecting it with one pole of an electric battery. Decom- 
position of the copper salt takes place, and copper is 
deposited upon the object. Alkaline solutions of the 
double cyanides are best adapted to the purpose. The 



[546 COLLEGE CHEMISTRY. 

process is extensively used in the preparation of electrotype 
plates. These are prepared either from wood-cuts or from 
type by making a mould of gutta-percha, covering this 
with graphite, and immersing the plate thus prepared in 
the copper-plating bath. The plate thus made is an exact 
reproduction of the wood-cut or type of which the impres- 
sion in gutta-percha was taken. 

Reactions which are of Special Value in Chemical 
Analysis. — Potassium or sodium hydroxide forms a blue 
precipitate which becomes black on standing or when 
heated. (See Oupric Hydroxide.) 

Ammonia first forms a greenish precipitate, which is a 
basic salt. With cupric sulphate the reaction takes place 
thus : 

S0 2 <2>Cu Cu<o H 

£ -f- 2NH 3 +2H 2 = X >S0 2 +(NH 4 ) 2 S0 4 . 

SO,<£>Cu Cll <OH 

If the action is carried farther, the basic salt dissolves, 
forming a dark blue solution. 

Potassium or sodium carbonate precipitates the basic 
carbonate referred to under Oupric Carbonate (which see). 
The change in color from blue to green which takes place 
in this precipitate is probably due to a loss of water. 

Potassium f err ocyanide, K 4 Fe(CN) 6 , forms a reddish- 
brown precipitate, which is the corresponding copper salt, 
Cu 2 Fe(ON") 6 . This compound is decomposed by caustic 
alkalies, forming cupric oxide and the corresponding 
alkali salt, JSTa 4 Fe(CN) 6 or K 4 Fe(CN) 6 . 

In the oxidizing flame the bead of borax or microcosmic 
salt is greenish blue, while when heated in the reducing 
flame it appears opaque and red. The red color is due to 
the reduction of the oxide to copper or cuprous oxide. 

Silver, Ag(At. Wt. 107.93). 

General. — In nearly all the compounds of silver the 
element is univalent, It, however, forms three oxides of 



SILVER. 547 

the formulas Ag 4 0, Ag 2 0, and AgO. The compounds 
correspond closely in many respects to the cuprous com- 
pounds. There is the same question here as in the case 
of copper as to whether the molecular weights correspond 
to the simple formulas AgCl, AgBr, AgN0 3 , etc., or to 
the doubled formulas Ag 2 Cl 2 , Ag 2 Br 2 , Ag 2 (N0 8 ) 2 , etc. 
There is no evidence at present available by which a 
decision between the two possibilities can be reached. 
The simpler formulas will therefore be used here. 

Forms in which Silver Occurs in Nature. — Silver occurs 
to some extent native, but for the most part in combina- 
tion, particularly with sulphur, and in company with lead. 
The principal ore of silver is the sulphide, Ag 2 S, which 
occurs in combination with other sulphides, as of lead, 
copper, arsenic, antimony, etc. The compounds with 
chlorine, bromine, and iodine are also found, but in 
smaller quantity than the sulphide. Small quantities of 
the sulphide are found in almost all varieties of galenite 
or lead sulphide. 

Metallurgy of Silver. — Much of the silver in use is 
obtained from galenite, PbS. This mineral is treated in 
such a way as to cause the separation of the lead (which 
see), and the silver is separated from sulphur at the same 
time. But it is dissolved in a large quantity of lead, and 
the problem which presents itself to the metallurgist is 
how to separate the small quantity of silver from the large 
quantity of lead. This is accomplished by melting the 
mixture and allowing it to cool until crystals appear. 
These are almost pure lead. They are dipped out by 
means of a sieve-like ladle, and the liquid left is again 
allowed to stand, when another crop of crystals is formed, 
and this can be removed in the same way as the first. By 
this means, and by again melting the crystals removed, 
allowing the liquid to crystallize, and removing the crystals 
formed, there is finally obtained a product which is rich 
in silver, but which still contains lead. This is heated in 
appropriate vessels in contact with the air, when the lead 
is oxidized, while the silver remains in the metallic state. 



548 COLLEGE CHEMISTRY. 

This is known as Pattinson's method. It has been super- 
seded by the 

Zinc Method, or Par Ices' s Method. — This consists in 
treating the molten alloy with zinc, which takes up all 
the silver, and the alloy of zinc and silver thus formed 
is removed, and afterwards treated with superheated 
steam, by which the zinc is oxidized while the silver is 
left unchanged. 

Amalgamation Process. — Some ores of silver are treated 
in another way, known as the amalgamation process. The 
ores are mixed with common salt and roasted, when the 
silver is obtained in the form of the chloride. This is then 
reduced to silver by means of iron and water, the reaction 
taking place as represented in the following equation : 

2AgCl + Fe = FeCl 2 + 2Ag. 

The mixture is next treated with mercury, which forms 
an amalgam with silver, while the other metals present do 
not combine with the mercury. The amalgam can be 
separated from the rest of the mass without much diffi- 
culty, and when heated to a sufficiently high temperature 
the mercury distils over, leaving the silver. 

Purification. — The silver in the market is not pure. 
For chemical purposes it can be purified by dissolving it 
in nitric acid, precipitating by means of hydrochloric 
acid, filtering and thoroughly washing the chloride, and 
reducing this either by melting it with sodium carbonate, 
or by pouring a little dilute hydrochloric acid upon it, 
and bringing a piece of zinc in contact with it, In the 
former case the reaction is 

2AgCl + Na 2 C0 3 = 2Ag + C0 2 + + 2NaCl; 

in the latter it is 

Zn + 2AgCl = ZnCl 2 + 2Ag. 

Properties. — Silver is a white metal with a high lustre, 
of specific gravity 10.5. It is not acted upon by the air, 
oxygen, or water. It melts at a lower temperature than 



SILVER. 549 

copper or gold, the melting-point being about 1000°. At 
the temperature of the oxyhydrogen blowpipe it distils, 
and in the experiments of Stas on the atomic weights of 
chlorine and silver the metal used was purified in this 
way. It is harder than gold and softer than copper, and 
its hardness is much increased by the addition of a little 
copper. It combines very readily with sulphur, forming 
black silver sulphide, and with chlorine, bromine, and 
iodine. The blackening of silver coins and other objects 
carried about the person is caused by the presence of 
minute quantities of sulphur compounds in the perspira- 
tion; and the blackening of spoons by contact with eggs 
is due to the presence of sulphur in the albumen of the 
eggs. When pure silver is melted in the air it absorbs 
about twenty times its volume of oxygen, and this is given 
off when the metal solidifies, causing in some cases a 
sputtering of the silver. This phenomenon is observed in 
the separation of silver from its ores in those processes in 
which it is necessary to melt the metal. It is known as 
"spitting." 

At the ordinary temperatures silver is converted into 
the peroxide, AgO, by ozone. When treated with hydro- 
chloric acid the metal becomes covered with a thin layer 
of the chloride, and no further action takes place, but it 
is dissolved easily by concentrated sulphuric acid and 
dilute nitric acid. With the concentrated acids reduction- 
products are formed as with copper. Silver is readily 
dissolved by a solution of potassium cyanide ; hence, such 
a solution is used in removing stains caused by silver salts. 
It is not acted upon by the alkaline hydroxides nor by 
potassium nitrate in the molten condition, while platinum 
is. Therefore silver vessels are used when it is desired 
to melt these substances in the laboratory, or to evaporate 
their solutions, as in the preparation of the caustic alkalies. 

Allotropic Forms, of Silver.— M. Carey Lea has dis- 
covered several curious allotropic forms of silver, the 
principal of which are briefly described by him as follows: 
"A. Soluble, deep red in solution, mat lilac, blue, or 



550 COLLEGE CHEMISTRY. 

green while moist, brilliant bluish-green metallic when 
dry. B. Insoluble, derived from A, dark reddish-brown 
while moist, when dry somewhat resembling A. C. Gold 
silver, dark brown while wet, when dry exactly resembling 
metallic gold in burnished lumps. Of this form there is 
a variety which is copper-colored. Insoluble in water; 
appears to have no corresponding soluble form." 

The form A is soluble in water, and the solution thus 
formed has a deep red color. The different varieties are 
formed by the action of reducing agents on solutions of 
silver salts. For example, the red soluble form is obtained 
by mixing dilute solutions of ferrous citrate and a silver 
salt. All the allotropic forms of silver are readily changed 
to the ordinary form. 

Lea further says: "All the forms of allotropic silver are 
sensitive to light. A when exposed to the sunlight soon 
becomes brown. The bright blue-green variety of B is 
changed into the pure gold-colored variety of C. Other 
forms of B turn brown on exposure to light." The red- 
yellow variety of C changes to bright gold color. " Con- 
tinued exposure seems to produce little further change so 
long as the substance is dry. But if the paper on which 
the silver is placed is kept moist by a wet pad, with three 
or four days of good sunshine, the change goes on nntil 
the silver becomes perfectly white and is apparently 
changed to normal silver." 

Alloys of Sliver. — For practical use, as in making coins 
and silver-ware, an alloy with copper is nsed, the pure 
metal being too soft. The alloy usually contains from 7-J 
to 10 per cent of copper. This alloy is harder than pure 
silver, and is capable of a higher polish. Silver amalgam 
is an alloy of silver and mercury, which is readily formed 
by bringing the two metals together. 

Argentous Chloride, Ag 2 Cl or Ag 4 Cl 3 , is formed by 
treating argentous salts with hydrochloric acid, and, 
possibly, to some extent when silver chloride, AgCl, is 
exposed to the light, though this is doubtful. 

Silver Chloride, Argentic Chloride, AgCl, is of special 



SILVER BROMIDE AND IODIDE. 55 x 

importance on account of its use in photography and in 
chemical analysis. It occurs to sonic extent in nature in 
Mexico and in the United States. It is easily formed as a 
white precipitate by adding hydrochloric acid to a solution 
of a silver salt, as, for example, the nitrate. In conse- 
quence of its insolubility in water it affords a convenient 
means of detecting silver and chlorine. If allowed to 
stand in the light it changes color, becoming first violet 
and finally black. This change in color appears to be due 
entirely to the reduction of the chloride to the. form of 
metallic silver. Concentrated hydrochloric acid dissolves 
it somewhat, and from this solution it crystallizes in octa- 
hedrons. An aqueous solution of ammonia dissolves it 
very easily, in consequence of the formation of a compound 
of the chloride with ammonia analogous to those formed 
by copper salts. The composition of the compound in the 
solution is, however, not known. 

Silver Bromide, AgBr, and Silver Iodide, Agl, are very 
similar to the chloride. Both occur in nature, and both 
are precipitated from solutions of silver salts by adding 
the corresponding nydrogen acids. The bromide is less 
easily soluble in ammonia than the chloride, and the 
iodide is almost insoluble in it. The bromide is formed 
by treating the chloride at the ordinary temperature with 
hydrobromic acid; and the iodide is formed from the 
chloride and from the bromide by treating these with 
hydriodic acid at ordinary temperatures. At higher tem- 
peratures, however, both the bromide and iodide are con- 
verted into the chloride by hydrochloric acid. Silver 
dissolves in concentrated hydriodic acid, and from the 
solution a salt of the formula AgT + HI or HAgI 2 is 
formed. It seems probable that this is a derivative of the 
acid H 2 I 2 , from which the double salt KI.AgI is also 
derived, as indicated in the formula KAgI 2 . Silver 
bromide at low temperatures is white, but easily changes 
to yellow, and by exposure it becomes darker, but not as 
readily as the chloride. The iodide is yellow, and under- 
goes change in the light only very slowly. The chloride 



55 2 COLLEGE CHEMISTRY. 

and iodide exist in several modifications, which differ from 
one another in their conduct towards light, and in their 
solubility. 

Application of the Chloride, Bromide, and Iodide of 
Silver in the Art of Photography. — The art of photography 
is based upon the changes which certain compounds, 
especially salts of silver, undergo when exposed to the 
light. Silver iodide is best adapted to most purposes. 
The salt is so changed by the light that when treated with 
certain compounds, such as ferrous sulphate, pyrogallic 
acid, etc., called "developers," a deposit of finely divided 
silver is formed upon the plate in those places affected by 
the light. A plate of glass or a sheet of properly prepared 
paper is covered in the dark with a thin layer of a salt of 
silver. The plate is then exposed in the camera to the 
action of the light which is reflected from the object to 
be photographed. According to the intensity of the light 
given off from the various parts of the object, the change 
of the silver salt takes place to a greater or less extent, and 
thus an image of the object is impressed upon the plate. 
But after the action of the "developer" is complete there 
is still upon the plate unchanged silver salt, and if this 
were now exposed to the light it would undergo change 
and the image would be obliterated. To remove this salt 
the plate is washed with a solution of sodium thiosulphate, 
Na 2 S 2 3 (hyposulphite), which dissolves the salt in conse- 
quence of the formation of a double salt of the formula 
2Na 2 S 2 3 .Ag 2 S 2 3 , which is readily soluble in water. 

Silver Triazoate, AgN 3 . — This is derived from triazoic 
acid (which see). It is formed by adding a solution of the 
acid to a solution of a silver salt. It is extremely explosive 
and should be dealt with very cautiously. Serious acci- 
dents have been caused by it. In appearance it resembles 
silver chloride, but it does not darken when exposed to the 
light. 

Silver Oxide, Ag 2 0. — The principal compound of silver 
and oxygen is that which has the composition Ag 2 0, and 



COMPOUNDS OF SILVER. 553 

in which the silver is univalent, as it is in its compounds 
with chlorine, bromine, and iodine. It is formed when a 
soluble hydroxide is added to a solution of a silver salt, 
and also by the action of concentrated solutions of the 
caustic alkalies on silver chloride. It is easily decomposed 
by heat and by reducing agents. 

Other Oxides of Silver. — Besides the ordinary oxide, 
silver forms a sub-oxide, Ag 4 0, corresponding to the sub- 
oxide of copper, Cu 4 0, and a peroxide of the formula AgO 
(or Ag 4 3 ), which is perhaps analogous to cupric oxide. 

Sulphides of Silver. — As has been stated, silver occurs 
in nature mostly in combination with sulphur as silver 
glance, Ag 2 S, which is in many minerals in combination 
with other sulphides. Examples of such double sulphides 
are the minerals stromeyerite, Cu 2 S. Ag 2 S, and pyrargyrite, 
3Ag 2 S.Sb 2 S 3 . 

Silver Nitrate, Argentic Nitrate, AgN0 3 . — This salt is 
formed by dissolving silver, or silver oxide, in nitric acid, 
evaporating to dryness, and heating until the salt is 
melted. It crystallizes in colorless rhombic plates. It is 
not changed in the light unless it comes in contact with 
organic substances, when it is reduced and metallic silver 
deposited. Hence the solution produces black spots on 
the fingers and clothing. As it melts easily, it is generally 
cast in small cylindrical moulds, and is found in the 
market in the form of thin sticks, and is known as lunar 
caustic. It disintegrates flesh, and is used in surgery as a 
caustic to remove superfluous growths. Owing to tie 
formation of a dark deposit when the salt is exposed to 
the light, it is used as a constituent of indelible inks. 

Silver Cyanide, AgCN, is formed as a caseous precipi- 
tate when a solution of hydrocyanic acid is added to a 
solution of silver nitrate. It does not change color in the 
light, is soluble in ammonia, but not in nitric acid. It 
readily forms double cyanides w T ith the cyanides of other 
metals. Of these, the salt with potassium cyanide, 
KAg(CN) 2 or KCN.AgCN, may be mentioned. 



554 COLLEGE CHEMISTRY. 

Reactions which are of Special Value in Chemical 
Analysis. — Hydrochloric acid precipitates insoluble silver 
chloride from solutions of silver salts, as silver nitrate. 

Soluble hydroxides precipitate silver oxide, not the 
hydroxide. Ammonia redissolves the precipitate in con- 
sequence of the formation of a compound of the oxide with 
ammonia of the composition Ag 2 0.2NH 3 . In dry condi- 
tion this salt is very explosive, and is known as fulminat- 
ing silver. 

Soluble carbonates precipitate the carbonate, Ag 2 C0 3 , 
which has a yellowish-white color. 

Ammonium carbonate redissolves the precipitate formed 
by it. 

Sodium phosphate, HNa 2 P0 4 , gives a precipitate of the 
normal salt Ag 3 P0 4 , which is yellow. 

Potassium ferrocyanide, K 4 Fe(CN) 6 , precipitates white 
silver ferrocyanide, Ag 4 Fe(CN) 6 . 

Potassium ferricyanide, K 3 Fe(CN) 6 , gives the corre- 
sponding silver salt, which is reddish brown. 

Potassium chromate or potassium bichromate (which see) 
gives a brownish-red precipitate of silver chromate. 

Gold, Au ( Vt. Wfc. 197.2). 

General. — Gold forms two series of compounds, in one 
of which it is univalent and in the other trivalent. In 
this respect it differs from the other members of the group. 
Examples of the compounds belonging to the two series 
are represented by the following formulas : 

AuCl AuCl 3 

Au Br AuBr 3 

Au 2 Au 2 3 

Those of the first series are called aurous compounds, 
those of the second series auric compounds. The basic 
character of gold is very weak, so that salts of the ordinary 
acids, as sulphuric, nitric, carbonic, etc., are not known. 
On the other hand, its higher oxide and hydroxide, 
Au(OH) 3 , have acid properties, and form salts similar in 



GOLD. 555 

composition to the meta-aluminates MA10 2 , and the 
metaborates MBO.,. These are the aurates, of which 
potassium aurate, KAuO, , is an example. So, also, the 
chloride combines readily with the chlorides of potassium 
and sodium, forming the clilor-aurates, KAuCl 4 , and 
NaAuCl 4 , which are analogous to the aurates. 

Forms in which Gold occurs in Nature. — Gold is gen- 
erally found in nature in the native condition — a fact 
which is due to the chemical inactivity of the elemert. 
That which is found in nature is never pure, but contains 
silver, and also, in different localities, iron, copper, and 
other metals. It is also found to some extent in com- 
bination with tellurium in the compounds AuTe 2 and 
(AuAg) 2 Te 3 . Native gold is frequently found enclosed in 
quartz, or more commonly in quartz sand. The principal 
localities in which it is found are California and some of 
the other Western United States, and Australia, Hungary, 
Siberia, and Africa. 

Metallurgy of Gold. — From the chemical point of view 
the metallurgy of gold is in general very simple. There 
are two kinds of gold mining — called placer mining and 
vein wining. In the former the earth and sand which 
contain gold are washed with water, which carries away 
the lighter particles, and leaves the gold mixed with other 
heavy materials. This mixture is then treated with mer- 
cury, which forms an amalgam with the gold, as it does 
with silver, and when this is placed in a properly con- 
structed retort and heated, the mercury passes over and 
leaves the gold behind. If silver is present, as is fre- 
quently the case, this is separated with the gold. In vein 
mining the gold ores are taken out of veins in the earth, 
and the gold separated by grinding the ores and treating 
them with mercury, as in the last stage of placer mining. 
Hydraulic mining is a modification of ordinary placer 
mining. It consists in forcing water under pressure 
against the sides of hills and mountains in which gold 
occurs loosely mixed with the earth. The earth is thus 
carried away and the heavier gold is deposited in sluices. 



55 6 COLLEGE CHEMISTRY. 

Some ores, especially those which contain tellurium, 
cannot be satisfactorily treated by the amalgamation 
process, and a method involving the use of potassium 
cyanide has been devised for them. In a solution of this 
salt gold dissolves, and from this solution it can be sep- 
arated in various ways. This method has come into 
extensive use of late years. 

Another process that is extensively used in the treat- 
ment of ores that do not give their gold to mercury is 
known as the chlorination process. This consists in treat- 
ing the finely ground ore with chlorine made from bleach- 
ing powder and sulphuric acid, and then precipitating 
the gold from the solution of the chloride by means of 
hydrogen sulphide. From the sulphide the metallic gold 
can be easily obtained. 

The gold obtained by any of the above methods is no" 
pure. It can be separated from silver by dissolving it in 
aqua regia, evaporating off the nitric acid, then diluting, 
and treating with a reducing agent, when metallic gold is 
precipitated. Thus when ferrous sulphate is used the 
following reaction takes place: 

3FeS0 4 + AuCl 3 = Fe 2 (S0 4 ), + Fe01 3 + Au. 

Another method of separating silver from an alloy with 
gold consists in treating the metal with nitric acid or with 
boiling concentrated sulphuric acid, which dissolves the 
silver and leaves the gold. This process is not satisfac- 
tory, however, unless the amount of gold in the alloy is 
less than 25 per cent. If the proportion of gold is greater 
than this, the alloy is melted with silver enough to bring 
the percentage of gold down to that mentioned. This is 
known as " quartation." 

Properties. — Gold is a yellow metal with a high lustre. 
It is quite soft, and extremely malleable, so that it is 
possible to make from it sheets the thickness of which is 
not more than 0.000002 millimeter. Thin sheets are 
translucent, and the transmitted light appears green. Its 
specific gravity is 19.3; its melting-point is higher than 



^■H 



ALLOYS, ETC., OF GOLD. 557 

that of copper, being about 1200°. It crystallizes in the 
regular system. Gold combines directly with chlorine, 
but not with oxygen. The three acids, hydrochloric, 
nitric, and sulphuric, do not act upon it; but aqua regia 
dissolves it, forming auric chloride, AuCl 3 , in consequence 
of the evolution of nascent chlorine. Molten caustic 
alkalies and their nitrates act upon it and form aurates. 

Alloys of Gold. — The principal alloy of gold is that 
which contains copper. The standard gold coin of the 
United States contains nine parts of gold to one of coj)per. 
The composition of gold used for jewelry is usually stated 
in terms of carats. Pure gold is 24-carat gold; 20-carat 
gold contains 20 parts of gold and 4 parts of copper; 
18-carat gold contains 18 parts of gold and 6 parts of 
copper, etc. Copper gives gold a reddish color, and makes 
it harder and more easily fusible. Gold is also alloyed 
with silver; and the alloy with mercury, known as gold- 
amalgam, is extensively used in the processes for extract- 
ing gold from its ores. 

Chlorides of Gold. — When gold is dissolved in aqua regia 
it is converted into auric chloride, AuCl 3 ; and if this solu- 
tion is evaporated a part of the chloride is decomposed 
into aurous chloride, AuCl, and chlorine. When gold is 
treated with dry chlorine it yields a mixture of auric 
chloride and metallic gold. When treated with a solution 
of stannous chloride a solution of auric chloride gives a 
purple-colored precipitate, known as the purple of Cassius, 
which appears to consist of finely- divided gold. 

Gold Sulphide, Au 2 S 2 . — This compound is precipitated 
together with sulphur from cold solutions of gold salts by 
means of hydrogen sulphide, and forms a brownish -black 
mass. It forms soluble compounds with the sulphides of 
the alkali metals. 

When hydrogen sulphide is passed into hot solutions of 
gold salts aurous -sulphide, Au 2 S, is thrown down as a 
steel-gray substance. This is soluble in pure water and is 
reprecipitated by hydrochloric acid. 



55 8 COLLEGE CHEMISTRY. 

EXPERIMENTS. 
Copper ais t d its Salts. 

Experiment 221.— Cuprous chloride is made by dissolving 
copper turnings in hydrochloric acid with the aid of a little nitric 
acid added from time to time. After solution is effected pour 
into water, when white cuprous chloride will be precipitated. 

Silver and its Salts. 

Experiment 222. — Dissolve a ten or twenty-five cent piece in 
dilute nitric acid. Dilute the solution to 200 to 300 cc. with hot 
water. Add a hot solution of common salt until it ceases to pro- 
duce a precipitate. Filter off the white silver chloride and wash 
with hot water. Dry the precipitate in the filter by placing the 
funnel with the filter and precipitate in an air-bath heated to 
100°-110°. Remove the precipitate from the filter and put it into a 
porcelain crucible. Heat gently with a small flame until the chlo- 
rine is melted ; then let it cool. Cut out a piece of sheet zinc large 
enough to cover the bottom of the crucible, and lay it on the silver 
chloride. Now add a little water and a few drops of dilute sul- 
phuric acid and let the whole stand for twenty-four hours. The 
silver chloride will be reduced to silver, and zinc chloride is 
formed : 

Zn + 2AgCl = ZnCU + 2Ag. 

Take out the piece of zinc and Wash the silver with a little dilute 
sulphuric acid, and then with water. Dissolve the silver in dilute 
nitric acid and evaporate to dryness in the water-bath, so that all 
the nitric acid is driven off. Dissolve the residue in water, and ; 
put the solution either in a bottle of dark glass or one wrapped; 
in dark paper. 

Experiment 223. — To a solution of silver nitrate containing 
about 5 grams of the salt in 100 cc. water, add a few drops of 
mercury, and let it stand. In a few days the silver will be depos- 
ited in the form of delicate crystals. The formation is called the 
"silver tree." 



CHAPTER XXIX. 

ELEMENTS OF FAMILY II, GROUP B: 
ZINC— CADMIUM.— MERCURY. 

General. — There is a strong resemblance between the 
first two elements of this group and magnesium, while 
mercury, in a general way, resembles the first two members 
of the copper group. Just as gold in the copper group 
furnishes a greater variety of compounds than the first 
two members of that group, so mercury furnishes a greater 
variety of compounds than the other members of the group 
to which it belongs. Zinc and cadmium, like magnesium, 
give only one class of compounds, and in these they are 
bivalent. The general formulas of some of the principal 
ones are : 

MCI,, M(OH) 2 , MO, MS0 4 , MC0 3 , MS. 

Mercury, on the other hand, furnishes two series of com- 
pounds, known as the mercurous and mercuric compounds, 
which correspond closely to the two series of copper salts. 
The power to form compounds belonging to both series is 
more strongly developed in mercury than in copper. 
Examples of the two classes are represented in the follow- 
ing formulas: 

Mercurous Compounds. Mercuric Compounds. 

HgCl HgCl 2 

Hgl Hgl, 

Hg 2 HgO 

HgNO," Hg(N0 3 ) 2 , etc. 

Just as the first member of Group A, Family II, beryl- 
lium, shows a somewhat acidic character in its hydroxide, 

559 



560 COLLEGE CHEMISTRY. 

while the other members of that group do not; so also the 
first member of Group B, Family II, zinc, is acidic, while 
the other members of the group are not. Glucinum 
hydroxide and zinc hydroxide dissolve in caustic alkalies, 
forming glucinates and zincates; while the hydroxides of 
all the other members of the two groups of this family are 
insoluble in caustic alkalies. 

Zinc, Zn (At. Wt. 65.4). 

General. — Zinc, in almost all its compounds, exhibits a 
close resemblance to magnesium. It always acts 'as a 
bivalent element. 

Forms in which it occurs in Nature. — Zinc occurs in 
nature in combination principally as the carbonate, or 
smithsonite, Zn00 3 ; as the sulphide, or sphalerite, ZnS; 
and as the silicate, Zn 2 Si0 4 . Among other compounds of 
zinc found in nature are gahnite, Zn(A10 2 ) 2 , and frank- 
linite, which contains the compound Zn(Fe0 2 ) 2 , with 
Fe(PeOJ 2 . 

Metallurgy. — The metallurgy of zinc is much simpler 
than that of magnesium, for the reason that the ores are 
easily converted into the oxide by roasting, and the oxide 
is easily reduced by heating it with charcoal. Owing to 
the volatility of the metal the vessels in which the reduc- 
tion is effected must be so constructed as to facilitate the 
condensation of the vapors. The vessels used are either 
earthenware muffles or tubes, open at one end and con- 
nected with iron receivers. At first the zinc vapor is con- 
densed in the form of a fine dust, as in the case of sulphur. 
This forms the commercial product called zinc dust. It 
always contains zinc oxide. Afterwards the zinc condenses 
to the form of a liquid, and this is cast in plates. The 
zinc thus obtained is not pure, but contains lead and iron, 
and sometimes arsenic and cadmium. It is called spelter. 
By repeated distillation it can be obtained pure. When 
distilled under diminished pressure, it is deposited in 
beautiful lustrous crystals, the forms of which are ex- 
tremely complicated. 



ZINC. 5 61 

Properties. — Zinc has a bluish-white color and a high 
lustre. The crystals above referred to, which are perfectly 
pure zinc, have a brilliant lustre, and do not change in the 
air. At different temperatures zinc has markedly different 
properties. At ordinary temperatures it is quite brittle; 
at 100°-150° it can be rolled out in sheets, but above 200° 
it becomes brittle again. It melts at 433°, and boils at 
1040°. When heated in the air it takes fire, and burns 
with a bluish flame, forming zinc oxide. This can be 
shown by means of the oxyhydrogen blowpipe. In dry 
air it does not change. Ordinary zinc dissolves in all the 
common acids, usually with evolution of hydrogen. In 
the case of nitric acid, however, the hydrogen acts upon 
the acid, reducing it to ammonia. The purer the zinc 
the less readily is it acted upon by sulphuric acid, and the 
pure crystals above referred to are scarcely acted upon at 
all by this acid. Zinc also dissolves in the caustic alkalies, 
forming zincates. Pure zinc can be made to act upon 
sulphuric acid by adding a few drops of platinum chloride. 

Applications. — Zinc is extensively used as sheet-zinc, in 
making galvanic batteries, for galvanizing iron, etc. Zinc 
dust is a very efficient reducing agent, either in alkaline 
or in acid solution. With caustic alkalies — as, for exam- 
ple, with potassium hydroxide — it gives hydrogen and a 
zincate : « 

Zn + 2KOH = Zn(OK) 2 + H,. 

With sulphuric acid also it gives hydrogen readily. Zinc 
is used in the preparation of important alloys. 

Alloys. — Iron covered with a layer of zinc is known as 
galvanized iron. As has been mentioned, zinc is a con- 
stituent of orass. It combines readily with mercury to 
form zinc amalgam, and this fact is taken advantage of 
for the purpose of preserving the zinc plates in galvanic 
batteries. Zinc plates covered with a layer of the amalgam 
are acted upon much more slowly than zinc itself. The 
amalgamation is effected by cleaning the zinc, dipping it 
in dilute sulphuric acid, and rubbing mercury over the 
surface with a brush or a piece of cloth. 



562 COLLEGE CHEMISTRY. 

Zinc Chloride, ZnCl 2 . — This is prepared by treating zinc 
with chlorine, or by dissolving zinc in hydrochloric acid, 
evaporating to dryness, and distilling the residue. It is a 
white deliquescent mass. From a very concentrated solu- 
tion in hydrochloric acid it is obtained in crystals of the 
composition ZnCl 2 -f- H 3 0. When the solution is evap- 
orated there is always some decomposition into basic 
chlorides, the hydroxide, and oxide. The basic chloride 
is formed thus : 

Zn<^J + HHO = Zn<^ H + HC1; 
the hydroxide thus: 

Zn< 01 H + HH0 = Zn< OH + HC1; 

and by higher heatiug the hydroxide yields the oxide and 
water : 

Zn<£|[ = ZnO + H.O. 

The chloride combines readily with water, and is used in 
the laboratory, as sulphuric acid and phosphorus pent- 
oxide are, for the purpose of extracting the elements of 
water from compounds. It has a caustic action, and is 
used in surgery on this account. Further, it acts as a 
disinfectant, and its solution is used for the purpose of 
preserving wood, particularly railroad sleepers, from decay. 

The chloride readily forms double chlorides like those 
formed by magnesium chloride. That formed with am- 
monium chloride, (NH 4 ) 2 ZnCl 4 , is made by mixing a 
solution of zinc in hydrochloric acid with a solution of 
ammonium chloride. This is used in soldering, as it 
cleans the surface of the metal, in consequence of the 
action of the zinc chloride on the oxides. 

Zinc Hydroxide, Zn(OH) 2 , is precipitated as a white 
amorphous powder when a soluble hydroxide is added to a 
solution of a zinc salt. It is redissolved by an excess of 
the reagent, and the zincate thus formed is decomposed 
on boiling, the hydroxide being reprecipitated. 



ZINC SULPHIDE. 565 

Zinc Oxide, ZnO, is formed in very finely divided con- 
dition by burning zinc in the air. The product is known 
as Flores zinci, and is sometimes called philosopher's wool. 
It is prepared on the large scale both by burning zinc and 
by heating the basic carbonate, which is formed by adding 
sodium carbonate to a solution of zinc sulphate. It is a 
white powder, that turns yellow when heated. Its chief 
use is as a constituent of paint under the name of zinc 
while. 

Zinc Sulphide, ZnS. — This compound occurs in nature, 
and is known as zinc blende. The mineral always contains 
a sulphide of iron, and also a small quantity of cadmium 
sulphide. When hydrogen sulphide is passed into a solu- 
tion of a zinc salt only a part of the zinc is thrown down 
as the sulphide, if the salt used is one of a strong acid, 
like sulphuric, nitric, or hydrochloric acid. The reason 
of this is that the sulphide is soluble in these acids, even 
when they are quite dilute. In the reaction the acid is set 
free and ionized, and although some sulphide is thrown 
down, the action soon stops: 

Zn + S0 4 + HH + S = ZnS + HH + S0 4 . 

If the acetate of zinc is used the precipitation is com- 
plete, because dilute acetic acid does not dissolve zinc 
sulphide, or because acetic acid is only slightly dissociated 
and therefore the disturbing hydrogen ions are not present. 
If sodium or potassium acetate is added to a solution of a 
neutral salt of zinc, hydrogen sulphide precipitates all the 
zinc, for the reason that the ions of the strong acid that 
are first formed react with the acetate and form the ions 
of the corresponding salt and also acetic acid, which exists 
in the solution largely in undissociated condition. Thus 
the hydrogen ions are eliminated, and the zinc sulphide is 
not dissolved : 

Zn + S0 4 + HH +~S = ZnS + HH + S0 4 ; 

gNa + 2CJELO, + HH + S0 4 = 2H 2 C 2 H 3 2 + 2Na + SO,. 



564 COLLEGE CHEMISTRY. 

The sulphide is, further, completely precipitated by 
soluble sulphides, as potassium and ammonium sulphides. 
Obtained by precipitation, zinc sulphide is a white amor- 
phous substance. 

Zinc Sulphate, ZnS0 4 . — This salt is readily formed by 
oxidation of the sulphide, and is hence found in nature 
accompanying the sulphide. It is manufactured by care- 
fully roasting zinc blende, and extracting with water. It 
crystallizes from the solution in water in large rhombic 
prisms of the composition ZnS0 4 -\- 7H 2 0. Like mag- 
nesium sulphate, it easily loses six molecules of water, but 
the last one is removed with difficulty. It appears, there- 
fore, that the constitution of the salt should be expressed 
by the formula (HO) 2 ZnS0 3 . Zinc sulphate, as has been 
stated (see p. 544), is commonly called white vitriol. It 
is easily reduced when heated with charcoal. The salt is 
used extensively in the preparation of cotton-prints and 
in medicine. 

Zinc Carbonate, ZnC0 3 , occurs in nature as smithsonite. 

Reactions which are of Special Value in Chemical 
Analysis. — The principal reactions which are made use of 
for the purpose of separating zinc from other elements 
have been mentioned above. These are the reactions with 
hydrogen sulphide and ammonium sulphide; with potas- 
sium and sodium hydroxide; with potassium and sodium 
carbonates; and with ammonium carbonate. Another 
reaction which is used in analysis is that which takes place 
when zinc salts are heated on charcoal before the blow- 
pipe, moistened with cobalt nitrate, and ignited. Under 
these circumstances a green-colored mass known as 
Einmann's green is formed, which is probably a zincate 

of cobalt, Zn < ^ > Co. 

Cadmium, Cd (At. Wt. 112.4). 

General. — The compounds of cadmium are similar to 
those of zinc and magnesium. The element occurs in 
nature in much smaller quantity than either of these, 



CADMIUM. 565 

often in company with zinc, and its compounds are not 
as frequently met with. It is always bivalent. A mineral 
known as greenockite is cadmium sulphide, CdS. 

Preparation and Properties. — Cadmium is obtained 
principally from different varieties of zinc blende, and 
separates with the zinc. Being more volatile than zinc, 
it passes over first when the mixture is distilled. From 
this first distillate, which contains the oxides of zinc and 
cadmium, the metals are reduced by heating with char- 
coal. It has a color like that of tin, and is harder than 
tin. According to the specific gravity of its vapor, its 
molecule is identical with its atom, for the molecular 
weight is approximately 112. 

Cadmium cldoride, CdCl 2 , like zinc chloride, is volatile; 
the sulphate crystallizes well, but is not analogous in com- 
position to the sulphates of magnesium and zinc, as the 
composition of the crystallized salt is represented by the 
formula 3CdS0 4 -f 8H 2 0; the normal carbonate, CdC0 3 , 
is precipitated by soluble carbonates. 

Cadmium Sulphide, CdS, is one of the most character- 
istic compounds of the element. It is a beautiful yellow 
substance, that is thrown down from a solution of a cad- 
mium salt by hydrogen sulphide. While it dissolves in 
concentrated acids it does not dissolve in dilute acids, and 
it is therefore completely precipitated by hydrogen sul- 
phide. It is used as a constituent of yellow paints. 

Analytical Reactions. — Cadmium, as has just been 
stated, is precipitated by hydrogen sulphide. It is thrown 
down together with the other elements of the hydrogen 
sulphide group (see p. 214). As the sulphide is not solu- 
ble in ammonium sulphide, it is easily separated from 
those of arsenic, antimony, and tin by treating with this 
reagent, when it is left undissolved in company with the 
sulphides of mercury, lead, bismuth, and copper. The 
double salt of cuprous cyanide and potassium cyanide is 
not decomposed by hydrogen sulphide, whereas the corre- 
sponding salt of cadmium is decomposed by it, and the 
yellow sulphide is precipitated. 



5 66 COLLEGE CHEMISTRY. 

The hydroxide of cadmium differs from that of zinc in 
not having acid properties. It does not dissolve in the 
caustic alkalies. 

Mercury, Hg (At. Wt. 200.3). 

General. — As already stated, mercury yields two series 
of compounds, known as mercurous and mercuric com- 
pounds, which are analogous to the two series of copper 
compounds. While, however, copper forms with the 
oxygen acids such salts as belong to the cupric series, as 
CuS0 4 , Cu(N0 3 ) 2 , etc., mercury forms salts belonging to 
both series. There is, for example, a mercurous nitrate, 
HgN0 3 , and a mercuric nitrate, Hg(N0 3 ) 2 ; a mercurous 
sulphate, Hg 2 S0 4 , and a mercuric sulphate, HgS0 4 , etc. 
The mercurous compounds are readily converted into the 
mercuric compounds by the action of oxidizing agents, 
and the mercuric are converted into mercurous compounds 
by the action of reducing agents. The action will be 
treated of under the individual compounds. The question 
as to the correct formula of the mercurous salts is in the 
same condition as that in regard to the formula of cuprous 
salts, with this difference : the molecular weight of mer- 
curous chloride leads to the formula HgCl, but there is 
evidence that when the chloride is heated some mercury is 
set free, and this has led to the suggestion that the mole- 
cule corresponds to the formula Hg 2 Cl 2 , and that the 
compound breaks down into mercury and mercuric chlo- 
ride when heated. It is, however, quite possible that the 
compound has the simpler formula, and that this under 
the influence of heat is partly decomposed, as represented 
in the equation 

2HgCl --= HgCl 2 + Hg. 

The fact that mercury is set free is, therefore, by no 
means satisfactory evidence that the formula of mercurous 
chloride is Hg 2 01 2 , and in the present state of the inquiry 
it is justifiable to write the formula HgCl. 

Forms in which Mercury occurs in Nature. — Mercury 
occurs native to some extent, but principally in the form 



MERCURY. 5 6 7 

of the sulphide, HgS, which is known as cinnabar. This 
is sometimes found crystallized, but generally amorphous. 
The chief localities are Idria, Almaden in Spain, and. New 
Almaden in California. 

Metallurgy of Mercury. — In order to obtain mercury 
from the sulphide this is roasted In vessels so constructed 
as to condense and collect the vapor of mercury given off. 
In the roasting process the sulphur is oxidized to sulphur 
dioxide, which oi course escapes. In some places the ore 
is mixed with limestone and distilled from clay or iron 
retorts, when the mercury passes over. Crude mercury is 
redistilled in order to purify it. It is also purified by 
treating it with dilute nitric acid or with a solution of 
ferric chloride. 

Properties. — Mercury is a silver-white metal of a high 
lustre. At ordinary temperatures it is liquid, though at 
— 39.5° it becomes solid. Its specific gravity is 13.5959. 
It does not change in the air at ordinary temperatures. 
It boils at 357.25°, and is converted into a colorless vapor, 
the specific gravity of which leads to the conclusion that, 
as in the case of cadmium, the molecule and atom are 
identical, or that the molecule consists of only one atom. 
It is insoluble in hydrochloric acid and in cold sulphuric 
acid; but dissolves in hot concentrated sulphuric acid, and 
is easily soluble in nitric acid. The vapor of mercury is 
very poisonous. 

Applications. — Mercury is extensively used in the manu- 
facture of thermometers, barometers, etc. ; as tin-amalgam 
for mirrors; and m the processes by which gold and silver 
are obtained from their ores. 

Amalgams. — The alloys which mercury forms with 
other metals are called amalgams. These compounds are 
generally obtained without difficulty simply by bringing 
mercury in contact with other metals. Among the amal- 
gams which are of- chief interest are those of sodium, 
ammonium, silver, and gold. Sodium amalgam is made 
by bringing mercury and sodium together. A crystallized 
amalgam containing the constituents m the proportions 



563 COLLEGE CHEMISTRY. 

represented in the formula Hg 6 Na has been obtained. 
Generally, sodium amalgam is easily decomposed by water, 
the mercury separating in the free state and the sodium 
acting upon the water, forming hydrogen and sodium 
hydroxide. It is much used in the laboratory as a con- 
venient means of producing hydrogen in alkaline solutions. 
It serves as an excellent reducing agent in some cases. 
Ammonium amalgam has already been referred to under 
the head of Ammonia (which see). It is a curious sub- 
stance, which is formed when an electric current acts upon 
a solution of ammonia containing some mercury which is 
connected with the negative pole, and also very easily by 
pouring a solution of ammonium chloride upon sodium 
amalgam. In the latter case sodium chloride and am- 
monium amalgam are formed. Apparently the reaction 
takes place according to the following equation : 

NH 4 C1 + NaHg = NaCl + NH 4 Hg. 

The product is extremely voluminous, and swells up 
during the reaction, so that it occupies under favorable 
circumstances about twenty times the volume occupied by 
the sodium amalgam. It has a metallic lustre, resembling 
in general the other amalgams. It is very unstable at the 
ordinary temperature, breaking down into mercury, hy- 
drogen, and ammonia. At a low temperature, however, 
it has been obtained in crystallized form. The metallic 
lustre and general outward appearance of the compound 
suggests that whatever is in combination with mercury in 
it has probably metallic properties, and this affords some 
confirmation of the ammonium theory, according to which 
the presence of the complex, NH 4 , in the salts formed by 
ammonia is assumed. Silver amalgam and gold amalgam 
vary in composition according to the method of prepara- 
tion, and when heated are comparatively easily decomposed. 
Mercurous Chloride, HgCl, is commonly called calomel. 
Like cuprous chloride, OuCl, and argentic chloride, AgOl, 
it is insoluble in water. It is formed most readily by 
reducing mercuric chloride. The reduction can be accom- 



MERCUR6US CHLORIDE-MERCURIC CHLORIDE. S () 

plished by means of Biilphurous acid, when the following 
reaction takes place: 

2HgCl a + 2H a + SO, = SHgCl + H a S0 4 + 2HC1. 

It is also formed by heating together mercuric chloride 
and mercury, and by subliming a mixture of mercuric 
sulphate, sodium chloride, and mercury. This method is 
the one mostly used in the manufacture of calomel. The 
product obtained by sublimation is crystalline; the precipi- 
tated substance forms a loose powder. As was stated 
above, the specific gravity of the vapor corresponds to that 
required for the formula HgCl. When acted upon for 
some time by light it undergoes partial decomposition into 
mercury and mercuric chloride. This is a fact of great 
importance, inasmuch as calomel is much used in medi- 
cine, and mercuric chloride is an active poison. Bottles 
in which calomel is kept should be carefully protected from 
the action of the light. 

Just as mercuric chloride is converted into mercurous 
chloride by reducing agents, so the latter is converted into 
the former by oxidizing agents. When, for example, 
mercurous chloride is treated with nitric acid it is con- 
verted into mercuric chloride and mercuric nitrate. 

Mercuric Chloride, or Corrosive Sublimate, HgCl 2 , is 
made by subliming a mixture of sodium chloride and mer- 
curic sulphate, 

HgS0 4 + 2NaCl = HgCl, + Na 2 S0 4 , 

and by dissolving mercury in aqua regia, evaporating to 
dryness, and subliming the residue. It is a white, trans- 
parent, crystalline mass; it is soluble in water, and can be 
obtained in crystalline form from the solution. It is more 
easily soluble in alcohol and ether than in water, and is 
extracted from a water solution by shaking with ether. 
It is quite volatile, and the specific gravity of its vapor 
corresponds to that required for the formula HgCl 2 . It 
is easily reduced to mercurous chloride by contact with 
organic substances, and by reducing agents in general. 



57o COLLEGE CHEMISTRY. 

The action of sulphur dioxide has already been treated of 
as furnishing a method for the preparation of mercurous 
chloride. Stannous chloride abstracts chlorine from it 
and forms mercurous chloride and metallic mercury, while 
the stannous chloride is converted into stannic chloride : 

2HgCl 2 + SnCl 2 = 2Hg01 + SnCl 4 ; 
HgCl 2 + SnCl 2 = Hg + SnCl 4 . 

Mercuric chloride is an active poison, and has been used 
extensively in this capacity. It has a very marked influ- 
ence upon the lower organisms, which play such an im- 
portant part in producing disease and the decay of organic 
substances, and is used as a disinfectant. Wood impreg- 
nated with a solution of it is partly protected from decay. 
In surgery it is used for the purpose of preventing 
contamination of wounds by the hands and instruments 
of the surgeon, it being customary now for the surgeon to 
wash his hands and instruments in a dilute solution of the 
chloride before performing an operation. 

Mercurous Iodide, Hgl, can be made by treating mer- 
cury with iodine; or by treating mercuric iodide with 
mercury; and more easily by adding potassium iodide to a 
solution of a mercurous salt, when it is thrown down as a 
greenish-yellow powder. It is unstable, and breaks down, 
yielding mercury and mercuric iodide. When treated 
with potassium iodide it suffers the same decomposition, 
the iodide which is formed combining with the potassium 
iodide, while mercury is deposited. Mercurous iodide is 
used in medicine. As it is more easily decomposed than 
mercurous chloride, and mercuric iodide is poisonous, care 
should be taken in its use. 

Mercuric Iodide, Hgl 2 , is made by direct combination 
of mercury and iodine, and by the action of potassium 
iodide upon a solution of a mercuric salt, when it is pre- 
cipitated as a beautiful scarlet-red powder. Though 
insoluble in water, it is soluble in alcohol and ether. It 
dissolves, further, in a solution of potassium iodide in 
consequence of the formation of a double iodide, or iodo- 



MERCURIC IODIDE- MERCURIC OXIDE. 571 

mercurate. When heated to about 150° the color changes 
from red to yellow, and after cooling the yellow substance 
changes to red. Sometimes it can be kept in the condition 
in which it has a yellow color for some time after it has 
cooled down, and then if touched at one point the change 
to the red substance takes place rapidly through the entire 
mass. Both the red and the yellow modifications are 
crystallized, but in different forms. The red crystals are 
tetragonal; the yellow ones are rhombic or monoclinic. 
The monoclinic structure is evidently, from what has been 
said, an unstable one for this compound. It seems prob- 
able that the difference between the two forms is due to a 
difference in the complexity of the molecules; perhaps the 
larger molecules of the red iodide are decomposed by heat 
into the smaller molecules of the yellow iodide. When 
the iodide is first precipitated from a solution of mercuric 
chloride by potassium iodide it is yellow, but it rapidly 
turns red. 

Mercuric Oxide, HgO, like the iodide, presents itself in 
two colors — the red and the yellow. When mercury is 
heated for a long time at a temperature near the boiling- 
point so that the air has access to it, it is converted into 
the red oxide. W T hen a solution of a mercuric salt is 
treated with caustic soda or caustic potash the yellow oxide 
is precipitated. W r hen heated, the red oxide becomes 
darker, and finally nearly black, and on cooling the red 
color reappears. The yellow oxide r,lso becomes darker 
on heating. Exposed to the light, the red oxide loses 
some of its oxygen, and mercury is deposited. When 
heated to a sufficiently high temperature both varieties 
give up all their oxygen, as was seen in the case of the red 
oxide in the preparation of oxygen. It will be remembered 
that oxygen was discovered by heating this compound. Of 
what importance this discovery was the student will now 
better appreciate than when the discovery was first men- 
tioned. It will now be seen that most of the chemical 
phenomena with which we have to deal involve the action 
of oxygen. 



$7 2 COLLEGE CHEMISTRY. 

Mercuric Sulphide, HgS, is the principal compound of 
mercury found in nature. This natural variety, which is 
known as cinnabar, is a red crystallized compound. When 
prepared by treating a solution of a mercuric salt with 
hydrogen sulphide it is an amorphous black powder. 
This black powder can, however, be converted into the 
red crystallized variety by sublimation, and by allowing it 
to stand in contact with a solution of the sulphide of an 
alkali metal. The sulphide, whether amorphous or crys- 
tallized, is acted upon with difficulty by acids. Dilute 
nitric acid, which easily dissolves the sulphides of those 
metals which resemble mercury, leaves it unacted upon, 
and advantage is taken of this fact for the purpose of 
separating it from the sulphides of lead, bismuth, copper, 
and cadmium in qualitative analysis. It is dissolved by 
concentrated nitric acid and by aqua regia. The substance 
used in medicine under the name Ethiops martialis is an 
intimate mixture of amorphous sulphide and sulphur. 
Like the oxide, cinnabar turns dark when heated, and it 
slowly undergoes the same change when exposed to the 
light, in consequence of a slight decomposition into mer- 
cury and sulphur. When that which has not been heated 
to too high a temperature is cooled down again, it acquires 
its original red color. If it has been heated to the tem- 
perature of sublimation, the black color of that which does 
not sublime is permanent. Cinnabar is used as a pigment. 

Mercuric Cyanide, Hg(CN) 2 , is formed by dissolving 
mercuric oxide in an aqueous solution of hydrocyanic acid. 
It is soluble in water, and crystallizes well in quadratic 
prisms. When heated it is decomposed into cyanogen and 
mercury, and this affords the most convenient method of 
making cyanogen. 

Fulminating mercury is an explosive compound much 
used in the manufacture of gun-caps. It is made by dis- 
solving mercury in nitric acid and adding alcohol. Its 
explosion consists in a sudden breaking down into nitro- 
gen, carbon dioxide, and mercury. The composition of 
the compound is represented by the formula C 3 N 2 2 Hg. 



AMMONIA COM POUNDS OF MERCURY. 573 

Mercurous Nitrate, HgN0 3 , is made by treating nitric 
acid with an excess of mercury. This salt is easily decom- 
posed, forming difficultly soluble basic salts, some of which 
are of complicated composition. The simplest is that 
which has the composition HgOH.HgN0 3 . 

Mercuric Nitrate, Hg(NO s ) 2 , is formed by treating 
mercuric oxide with an excess of nitric acid and evaporat- 
ing, when the salt can, under favorable circumstances, be 
obtained in crystals. It is easily decomposed by water, 
with formation of a basic salt which is insoluble in water. 

Compounds formed by Salts of Mercury with Ammonia. 
— "When mercuric chloride is treated with ammonia a white 
precipitate is formed, the composition of which is repre- 
sented by the formula HgClNII,. The formation takes 
place according to the equation 

HgCl 2 + 2XH 3 = HgClNH, + NH 4 C1. 

The compound is known as white precipitate, or, to distin- 
guish it from another similar compound, as infusible white 
precipitate. The latter is formed by adding a solution of 
mercuric chloride to a boiling solution of ammonium 
chloride containing ammonia, and separates out on cooling. 
It has the composition Hg(NH s Cl) 2 . This is called mer- 
curic diammonium chloride, or fusible white precipitate. 
Mercurous compounds corresponding to both the above- 
mentioned derivatives of mercuric chloride are known. 
The first, or mercurous chlor amide, Hg 2 NH 2 Cl, is the black 
substance formed when mercurous chloride is treated with 
ammonia: 

2HgCl + 2NH 3 = Hg 2 NH 2 Cl + NH 4 C1. 

The second, or mercurous ammonium chloride, HgNH 3 Cl, 
is formed by treating calomel with ammonia gas. 

Similar compounds are obtained from the other salts of 
mercury. In a solution of mercurous nitrate ammonia 
forms a black precipitate, known in pharmacy as Mercurius 
solubilis Hahnemanni. 

Reactions which are of Special Value in Chemical 
Analysis. — As has been seen in the account already given 



574 COLLEGE CHEMISTRY. 

of the conduct of the compounds of mercury, mercurous 
and mercuric compounds conduct themselves quite differ- 
ently. Among the most characteristic reactions of mer- 
curous compounds are the following: 

Sodium or potassium hydroxide gives a black precipitate 
of mercurous oxide. 

Ammonia gives a black precipitate which is a compound 
of mercurous oxide and ammonia. 

Hydrochloric acid and soluble chlorides form mercurous 
chloride, and the precipitate turns black when treated with 
ammonia. 

The reactions with stannous chloride, potassium iodide, 
hydrogen sulphide, and ammonium sulphide have been 
explained. 

The principal reactions of the mercuric salts are the 
following: Sodium or potassium hydroxide produces a 
yellow precipitate of mercuric oxide. 

The reactions with hydrogen sulphide, stannous chlo- 
ride, and potassium iodide have been described above. 

ELEMENTS OF FAMILY III, GROUP B: 
GALLIUM.— INDIUM.— THALLIUM. 

General. — All the elements of this group are rare. 
Gallium forms compounds which in composition are 
analogous to those of aluminium, in which it is trivalent, 
as in the compounds GaCl 3 , Ga 2 3 , Ga(N0 3 ) 3 , Ga 2 (S0 4 ) 3 , 
etc. On the other hand, it also forms compounds in 
which it is bivalent, as in the chloride GaCl 2 and GaO. 
Indium also forms compounds in which it is trivalent, and 
a few in which it is bivalent. Thallium, like gold, is tri- 
valent and univalent in its compounds. 

Gallium, Ga (At. Wt, 70).— This element is found in 
some varieties of zinc blende, and was discovered in that 
which occurs at Pierrefitte, in France. It owes its name 
to the Latin name of France, Gallia. Like scandium, it 
is of special interest for the reason that its properties were 
foretold by Mendeleeff some years before it was discovered, 
and it was described by him under the name of eka- 



INDIUM— THALLIUM. 575 

aluminium, as scandium was described under the name of 
eka-boron. 

Indium, In (At. AVt. 114). — Indium was discovered in 
a variety of zinc blende found at Freiberg, in Germany. 
The discovery was made by means of the spectroscope, 
and as the spectrum of the element contains two very 
characteristic bine lines it was called indium. It has 
since been found in other varieties of zinc blende, but 
always in small quantity. 

Thallium, Tl (At, Wt. 20-1.1).— This element was dis- 
covered in the flue-dust of a sulphuric-acid factory in the 
Harz mountains, by the aid of the spectroscope. As it 
colors the flame a beautiful green it was called thallium, 
from the Greek SaXXo?, which signifies a green branch. 
It has since been found in a number of varieties of iron 
pyrites and copper pyrites. 

EXPERIMENTS. 

Zinc and its Salts. 

Experiment 224. — Heat a small piece of zinc on charcoal in 
the oxidizing flame of the blowpipe. The white fumes of zinc 
oxide (philosopher's wool) will be seen, and the charcoal will be 
covered with a film which is yellow while hot, but becomes white 
on cooling. 

Experiment 225. — Dissolve some zinc dust in a solution of 
sodium hydroxide, and see whether hydrogen is given off. 

Mercury and its Salts. 

Experiment 226. — Make a solution of mercurous nitrate by 
treating at the ordinary temperature an excess of mercury with 
nitric acid, which is not too concentrated ; and with this solution 
study the conduct of mercurous salts. 

Experiment 227. — Heat some of the solution of mercurous 
nitrate to boiling, then add a few drops of concentrated nitric 
acid, and boil again. With the solution thus obtained study the 
conduct of mercuric salts. 



CHAPTER XXX. 

ELEMENTS OF FAMILY IV, GROUP B; 
GERMANIUM.— TIN.— LEAD. 

General. — Of the elements of this group germanium is 
extremely rare. It was discovered by Winkler in 1885. 
Tin and lead, on the other hand, have long been known, 
and are extensively used. All form two series of com- 
pounds, in one of which they are bivalent and in the other 
quadrivalent. The general formulas of some of the prin- 
cipal compounds of the first series are as follows : 

M01 2 , MO, M(OH) 2 , M(N0 3 ) 2 , MS0 4 , MC0 3 , etc. 

The general formulas of some of the principal compounds 
of the other series are as follows : 

MC1 4 , M0 2 , M(OH) 4 , MO(OH) 2 , etc. 

These elements have already been referred to at the close 
of Chapter XXII (see p. 399), and attention was there 
called to the resemblance between them and carbon and 
silicon. In this connection it will be well to repeat what 
was there said. Of the three elements of the group, 
germanium and tin are more acidic in character than lead. 
They combine with chlorine in two proportions, forming 
the chlorides GeCl 2 , SnCl 2 , PbCl 2 , GeCl 4 , SnCl 4 , PbCl r 
With oxygen they combine to form the compounds Ge0 2 , 
Sn0 2 , and Pb0 2 . Stannic oxide, Sn0 2 , and lead peroxide, 
Pb0 2 , form salts with bases, and these have the composi- 
tion represented by the general formulas M 2 Sn0 3 and 
M ? Pb0 3 , and are therefore analogous to the silicates, car- 

576 



GERMANIUM- TIN. 577 

bonates, and titanates. On the other hand, further, salts 
are known which are derived from the oxide PbO. These 
have the general formula M 2 Pb0 2 , and are to be regarded 
as salts of an acid Pb(OH) 2 . These salts are not stable, 
and are not easily obtained. Most of the derivatives of 
lead are those in which it plays the part of a base-forming 
element. Notwithstanding the marked analogy between 
some of the compounds of tin and those of the members 
of the silicon group, it appears, on the whole, advisable to 
treat of this element in company with lead, which it also 
resembles in many respects. 

Germanium, Ge (At. Wt. 72). 

Germanium is the third element the properties of which 
were foretold by Mendeleeff by the aid of the periodic law. 
As it occurs in the silicon group he called it eka-silicon. 
It was discovered in a silver ore occurring at Freiberg, 
Germany. The name has, of course, reference to the 
country in which it was discovered. It acts mostly as a 
base-forming element, being perhaps more like tin than 
any other one metal. It forms the two chlorides GeCl 2 
and GeCl 4 , and the corresponding fluorides GeF 2 and 
GeF 4 ; but preferably it forms those compounds in which 
the element is quadrivalent. The fluoride forms double 
salts resembling the fluosilicates, as, for example, potas- 
sium fluogermanate, K 2 GeF 6 . 

Tin", Sn (At. Wt. 118.5). 

General. — The compounds of tin with which we gen- 
erally have to deal belong, with the exception of stannous 
chloride, to the series in which the metal is quadrivalent, 
and in this series it acts as an acid-forming element. The 
chloride, SnCl 4 , corresponds to the chlorides of carbon 
and silicon, CC1 4 and SiCl 4 . Unlike these elements, how- 
ever, it does not form a compound with hydrogen. 

Occurrence. — Tin occurs almost exclusively as tin-stone 
or cassiterite in nature. This is the dioxide, Sn0 2 , corre- 



57^ COLLEGE CHEMISTRY, 

sponding to carbon dioxide, 00 2 : silicon dioxide, Si0 2 ; 
titanium dioxide, Ti0 2 ; etc. It also occurs in small 
quantities in company with gold as metallic tin, and in a 
variety of pyrites of the formula Cu 4 SnS 4 -J- FeSnS 4 , 
known as stannite. 

Metallurgy. — The ores are roasted for the purpose of 
getting rid of the sulphur and arsenic, and the oxide is 
then heated with coal in a furnace. After the reduction 
is complete the tin is drawn off and cast in bars. This 
tin is impure, and when again slowly melted, that which 
first melts is purer. By letting it run off as soon as it 
melts, the comparatively difficultly fusible alloy remains 
behind, and the tin is thus rendered much purer. The 
commercial variety of tin known as Banca tin is the purest. 
It receives its name from Banca, in the East Indies, where 
it is made. Block-tin is made in England, and is also 
comparatively pure. 

Properties. — Tin is a white metal, which in general 
appearance resembles silver. It is soft and malleable, and 
can be hammered out into very thin sheets, forming the 
well-known tin -foil. Its specific gravity is 7.3. At 200° 
it is brittle, and at 228° it melts. At ordinary tempera- 
tures it remains unchanged in the air. It dissolves in 
hydrochloric acid, forming stannous chloride, Sn01 2 ; in 
sulphuric acid, forming stannous sulphate, SnS0 4 , sulphur 
dioxide being evolved at the same time. Ordinary con- 
centrated nitric acid oxidizes it; the product being a 
compound of tin, oxygen, and hydrogen, known as meta- 
stannic acid, which is a white powder insoluble in' nitric 
acid and in water. It is dissolved by a hot solution of 
potassium hydroxide which forms potassium stannate, 
K 2 Sn0 3 . 

Applications. — It is used in making alloys, of which 
bronze (see p. 540), soft solder, and britannia metal are 
the most important. It is used also for protecting other 
metals, as in the tinware vessels in such common use, 
which are made of iron covered with a layer of tin, 
Qopper vessels are also frequently covered with tin. 



STANNOUS CHLORIDE. 579 

Alloys. — Bronze has already been treated of under 
Copper. It is made of copper, tin, and zinc. Soft solder 
is made of equal parts of tin and lead, or of two parts of 
tin and one of lead. Britannia metal is composed of nine 
parts of tin and one of antimony. Tin amalgam is made 
by bringing tin and mercury together, and is used in the 
silvering of mirrors. Alloys of tin and antimony are much 
used for machine bearings. Those in most common use 
contain not only these two elements, but also copper and 
lead. 

Stannous Chloride, SnCl 2 , is formed by dissolving tin 
in hydrochloric acid, and if the solution is concentrated 
enough the compound crystallizes out. The crystals have 
the composition SnCl 2 4- 2H 2 0. This is the commercial 
product known as tin salt. It is easily soluble in water, 
but if the solution is dilute it becomes turbid in conse- 
quence of the formation of the insoluble basic salt, 

2Sn<£ 1 H +H 2 0: 

2SnCl 2 + 3H 2 = 2Sn<^ H + H 2 + 2H01. 

This same decomposition takes place when the solution is 
allowed to stand in contact with the air. Under these 
circumstances a part of the stannous chloride is converted 
into stannic chloride by oxidation: 

3SnCl 2 + H 2 + = SnCl 4 + 2Sn(OH)Cl. 

When the crystals are heated they melt at about 40°, and 
if heated higher they undergo partial decomposition, 
forming the oxide and hydrochloric acid : 

SnCl 2 + H 2 = SnO +2HC1. 

Stannous chloride has a marked tendency to combine 
with chlorine and to pass over into stannic chloride. This 
power has already been shown in its action upon mercuric 
chloride and. upon mercurous chloride, It reduces the 



580 COLLEGE CHEMISTRY. 

former, first to mercurous chloride, and it then abstracts 
the chlorine from this, leaving metallic mercury: 

2HgCl 2 + SnCl 3 = 2HgCl + SnCl 4 ; 
2Hg01 + Sn01 2 = 2Hg + SnCl 4 . 

Stannous chloride is a good mordant. 

Stannic Chloride, SnCl 4 , formed by treating tin with 
chlorine, is a colorless liquid which boils at 114°, and the 
specific gravity of its vapor is that required by a compound 
of the formula SnCl 4 . In the air it gives off fumes in 
consequence of the action of moisture. It has long been 
known by the name spiritus fumans Libavii. It combines 
with water, forming a number of crystallized hydrates. 
When its solution in water is boiled stannic acid is pre- 
cipitated : 

SnCl 4 + 3H 2 = H 2 Sn0 3 + 4HC1. 

Stannic chloride is used as a mordant. With other 
chlorides it forms the chlorost annates of the general 
formula M 2 SnCl 6 or SnCl 4 .2MCl, of which the ammonium 
salt, (NH 4 ) 2 SnCl 6 , or pink salt, is the best known. This 
is manufactured for use in making cotton-prints. 

Stannic Hydroxide, Sn(OH) 4 , is perhaps formed when 
a solution of stannic chloride in water is boiled. The 
precipitate obtained has, however, the composition H 2 Sn0 3 , 
and this is known as stannic acid. Stannic acid is precipi- 
tated also by treating a solution of a stannate with just 
enough acid to effect decomposition. The decomposition 
with hydrochloric acid takes place as represented in the 
equation 

Na 2 Sn0 3 + 2HC1 = 2NaCl + H 2 Sn0 3 . 

The compound thus obtained is insoluble in water, but is 
easily soluble in hydrochloric, nitric, and sulphuric acids, 
and in the caustic alkalies. With the alkalies it forms 
stannates, as sodium stannate, Na 2 Sn0 3 , and potassium 



STANNIC OXIDE, ETC. 581 

stan?iate, K,Sn(>... The former is made on the large scale, 
and is known as preparing-salt. 

Metastannic Acid. — When tin is treated with concen- 
trated nitric acid it i& converted into a white powder which 
is insoluble in water and in acids, but nevertheless seems 
to be a hydroxide of tin of the same composition as stannic 
acid. This is kiipwn as metastannic acid. With alkalies 
it forms salts which in properties and composition are 
entirely different from the stannates. They are known as 
the metastan nates. Two sodium salts are known, which 
differ in composition. One of these has the composition 
Na 2 Sn 5 O u , the other is Na 2 Sn 9 19 . They are probably 
derived from acids analogous to the polysilicic acids, which 
may be called poly stannic acids. The question as to the 
composition of metastannic acid is an open one. 

Stannic Oxide, Sn0 2 , as has been stated, is the principal 
form in which tin occurs in nature. The mineral is known 
as cassiterite or tin-stone. It is found at Cornwall, 
England; on the East Indian islands Banca and Biliton; 
and in Malacca. It crystallizes in the tetragonal system, 
and is generally colored from brown to lack. It is formed 
as a white powder by burning tin in the air, and by heat- 
ing the different varieties of stannic hydroxide. It is 
infusible, and is not acted upon by concentrated hydro- 
chloric or nitric acid. Concentrated sulphuric acid, how- 
ever, converts it into a gelatinous liquid from which 
stannic oxide is precipitated by water. 

Stannous Sulphide, SnS, which is formed when hy- 
drogen sulphide is passed into a solution of stannous 
chloride, is a brownish-black powder. When treated with 
the soluble sulphides and sulphur, or with a soluble poly- 
sulphide, it dissolves, forming a sulphostannate as repre- 
sented in these equations: 

SnS + (NH 4 ) 2 S + S = (NH^SnS,; 
SnS +^(NH 4 ) 2 S 2 = (NH 4 ) 2 SnS 3 . 

Stannic Sulphide, SnS 2 . — This compound is obtained 
in crystallized form by heating together tin-filings, sulphur, 



S 82 COLLEGE CHEMISTRY. 

and dry ammonium chloride; and in amorphous form by 
treating a solution of stannic chloride with hydrogen sul- 
phide. In the former case it is a golden-yellow crystalline 
substance; in the latter a yellow powder. The crystalline 
variety is known as mosaic gold. When heated to a high 
temperature it is converted into stannous sulphide. The 
precipitated variety is easily dissolved i»y concentrated 
hydrochloric acid, and converted into metastaunic acid by 
concentrated nitric acid. The crystallized variety is not 
soluble in hydrochloric acid, and is affected but slightly 
by nitric acid. The crystallized variety, or mosaic gold, 
is used as a pigment, particularly for bronzing. Stannic 
sulphide dissolves easily in the soluble sulphides, forming 
sulphostannates : 

SnS 2 + (OT 4 ) 2 S = (NH 4 ) 2 SnS 3 ; 
SnS 2 + K 2 S = K 2 SnS 3 . 

Reactions which are of Special Value in Chemical 
Analysis. — The reactions with hydrogen sulphide, and the 
conduct of the precipitated sulphides when treated with 
ammonium sulphide or polysulphide, are constantly 
utilized for analytical purposes when tin is present. 

The reducing action of stannous compounds serves to 
distinguish them from stannic compounds. 

The carbonates of the alkali metals precipitate the 
hydroxides, and these do not dissolve in an excess of the 
carbonate. Metallic zinc precipitates the tin from a solu- 
tion of a tin compound as a spongy mass, If the precipi- 
tation is carried on in a platinum vessel, the platinum is 
not colored by it. By this means it can be distinguished 
from antimony, which is reduced under the same circum- 
stances, but is deposited as a black coating upon the 
platinum. 

Lead, Pb (At, Wt. 206.9). 

General. — The basic properties of lead are stronger, and 
its acidic properties weaker, than those of tin. Its prin- 



LEAD. 5 8 3 

cipaJ compounds are those in which it acts as a base-form- 
ing element. The compounds in which it is quadrivalent, 
such as Pb0 2 , are comparatively unstable, and when treated 
with acids they readily pass over into the compounds of 
the series in which the lead is bivalent. Thus lead oxide 
itself readily gives up oxygen when treated with acids, and 
yields salts of bivalent lead. 

Forms in which Lead occurs in Nature. — Lead occurs 
in nature as the sulphide, PbS, which is known as galenite 
or galena. Other natural compounds of the metal are the 
carbonate, PbC0 3 , known as cerussite; the phosphate, 
Pb 3 (P0 4 ) 2 ; the chromate, PbCrO t , or crocoisite; and the 
molybdate, PbMo0 4 , or wulfenite. 

Metallurgy. — Most of the lead in the market is obtained 
from the sulphide, and as most of the sulphide contains 
silver both metals have to be considered in the treatment 
of the ore. Under the head of Silver (which see) reference 
was made to the methods by which this metal is separated 
from the lead after both have been separated from their 
compounds. Here it will only be necessary to show how 
the metals are extracted together from the ore. This is 
accomplished in one of two ways : 

(1) By heating the sulphide with iron, when the latter 
combines with the sulphur, forming iron sulphide, while 
the lead is set free. 

(2) By roasting the sulphide until it is partly converted 
into lead oxide and lead sulphate, and then heating the 
mixture without access of air, when two reactions take 
place, which are represented in these equations : 

PbS + 2PbO = 3Pb + S0 2 ; 
PbS + PbS0 4 = 2Pb + 2S0 2 . 

The lead is thus set free, and the sulphur is driven off as 
sulphur dioxide. 

Properties. — Lead is a bluish-gray metal, with a high 
lustre. It is very soft, and not very strong; melts at 
325°, and has the specific gravity 11.37. At a high tem- 
perature it is converted into vapor. When heated in con- 



5^4 COLLEGE CHEMISTRY. 

tact with the air it becomes covered with a layer of oxide, 
as can be seen in a vessel containing the molten metal. 
Pare water acts upon lead when air has access to it, and 
some of the lead dissolves. If the water contains salts in 
solution, such as calcium carbonate, gypsum, etc., or if it 
contains carbon dioxide, it acts only very slightly upon the 
metal; further, water that contains organic matter in a 
state of decomposition dissolves lead with comparative 
ease. Nitric acid dissolves lead readily; as, however, lead 
nitrate is insoluble in nitric acid, it is necessary to use 
comparatively dilute acid. Concentrated sulphuric acid 
dissolves lead to some extent, and therefore a little lead 
sulphate is sometimes contained in commercial sulphuric 
acid. When such a solution is diluted with water the 
sulphate is precipitated. Hence when some specimens of 
the commercial acid are diluted with water they become 
turbid, and after standing for a time they become clear, 
as the lead sulphate settles. Hydrochloric acid acts only 
slightly upon lead. Acetic acid dissolves the metal very 
readily. It is precipitated in metallic form from a solu- 
tion of one of its salts by metallic zinc. The formation is 
sometimes very beautiful, especially if the zinc is sus- 
pended in the solution. It is called the "lead tree," or 
Arbor Satumi. The action consists in a substitution of 
zinc for the lead. After the action is complete all the 
lead is deposited as metallic lead, and the zinc has entered 
into its place, forming a salt which remains in solution. 
Thus, if lead nitrate is used, zinc nitrate is in the solution. 
Applications. — Lead is extensively used for a variety of 
purposes, as, for example, for making sulphuric-acid 
chambers, for evaporating-pans for alum and sulphuric 
acid, for shot, for water-pipes, and in making various 
alloys. The use of lead water-pipes is a matter of much 
importance from the sanitary point of view, as is evident 
from the statements above made concerning the action of 
water upon the metal. Ordinary drinking-water acts 
under most circumstances only very slightly upon lead, 
and not enough is dissolved to be dangerous to those using 



LEAD CHLORIDE, ETC. 5 8 5 

fche water. At the same time circumstances may at any 
time arise which will increase the solvent power of the 
water, and thus cause serious results; and it would un- 
doubtedly be better if the use of such pipes could be 
entirely avoided in cases in which the water is to be used 
for drinking purposes. Babbitt's metal, which is used for 
machine bearings, contains copper, antimony, tin, and 
lead. 

Lead Chloride, PbCl 2 , is formed when hydrochloric acid 
or a soluble chloride is added to a cold solution of a lead 
salt, and appears as a white precipitate. It is soluble in 
hot water, and is deposited in the form of long, needle- 
shaped crystals when the solution cools. It occurs in 
nature in small quantity, as cotunnite. 

Lead Tetrachloride, PbCl 4 , is a liquid that forms a 
crystalline mass at — 15°. It is decomposed into the 
dichloride and chlorine at the ordinary temperature. At 
105° this takes place with explosion. It is formed by 
passing chlorine into concentrated hydrochloric acid con- 
taining lead chloride. After saturation, ammonium chlo- 
ride is added, when the double salt, PbCl 4 .2NH 4 Cl, 
separates in yellow crsytals. Ice-cold concentrated sul- 
phuric acid decomposes it into ammonium sulphate, 
hydrochloric acid, and lead tetrachloride, which collects 
as a yellow oil below the sulphuric acid. 

Lead Iodide, Pbl 2 , is a yellow substance that crystallizes 
from water in beautiful lustrous laminae. It is precipi- 
tated when potassium iodide is added to a solution of a 
lead salt. 

Oxides of Lead. — Lead forms five distinct compounds 
with oxygen, the formulas and names of which are as fol- 
lows: lead suboxide, Pb 2 0; lead oxide, PbO; lead ses- 
quioxide, Pb 2 3 ; minium, Pb 3 4 ; and lead peroxide, Pb0 2 . 
Lead Suboxide, Pb 2 0. — This compound is formed when 
lead is exposed to the air. 

Lead Oxide, PbO. — This is formed by heating lead 
nitrate, and is then left behind in the form of a yellow 
powder. If heated to fusion it solidifies, forming a 



5^6 COLLEGE CHEMISTRY. 

yellowish or reddish mass known as litharge. This is 
formed in large quantity in the process of separating silver 
from lead. With the strongest bases it forms salts similar 
to those formed by zinc. This is seen in the solubility of 
the hydroxide in sodium and potassium hydroxide, which 
is due to the formation of compounds known as plumbites : 

Pb(OH) 2 + 2KOH = Pb(OK) 2 + 2H 2 0. 

Heated with silicon dioxide it forms a silicate which is 
easily fusible. Lead oxide is used extensively in the 
manufacture of flint glass for optical purposes, as was 
described under Glass (which see). It also finds applica- 
tion in glass painting and porcelain painting, and is used 
in the manufacture of lead salts, particularly " sugar of 
lead," which is an acetate. 

Lead Peroxide, Pb0 2 , is formed by treating minium or 
red lead with dilute nitric acid. Minium has the compo- 
sition Pb 3 4 . When treated with nitric acid, a part dis- 
solves as lead nitrate, and lead peroxide remains behind, 
as represented in the equation 

Pb 3 4 + 4HN0 3 = Pb0 2 + 2Pb(NO s ) 2 + 2H 2 0. 

The peroxide is formed in general by the action of oxidiz- 
ing agents upon the lower oxides of lead. One of the 
most convenient methods of making it consists in treating 
lead acetate with a filtered solution of bleaching-powder. 
It is a dark-brown powder, insoluble in water. When 
ignited it loses half of its oxygen, and it gives up its 
oxygen readily to other substances. Towards hydrochloric 
acid it acts like manganese dioxide, giving lead chloride 
and chlorine according to the equation 

Pb0 2 + 4HC1 = PbCl 2 + 2H 2 + Cl 2 . 

It appears probable that the tetrachloride is first formed, 
and that this then breaks down into the dichloride and 
chloride. When the peroxide is treated in the cold with 
hydrochloric acid it dissolves, and when this solution is 



STORAGE BATTERY. 5S7 

heated it gives off chlorine. Farther, when it is treated 

with caustic alkalies lead peroxide is thrown down. 

Lead peroxide dissolves in concentrated caustic potash 

and forms a salt of the formula K 2 Pb0 3 , or PbO< r>K , 

analogous to potassium stannate, K 2 Sn0 3 , silicate, K 2 SiO s , 
carbonate, K 2 C0 3 , etc. Other salts derived from the acid 
PbO(Oll)., are known, and are called plumbates. 

Storage Battery. — An efficient galvanic battery can be 
made by putting two lead plates in dilute sulphuric acid, 
one of the plates being covered with a layer of lead 
peroxide. The sulphuric acid acts upon the lead, convert- 
ing it into lead sulphate, which forms a solid layer upon 
the plate. On the other hand, the acid acts upon the 
lead peroxide, reducing the lead to the bivalent condition 
and forming with it lead sulphate. As long as the chem- 
ical action can take place, that is to say, as long as there is 
lead peroxide present, the battery continues to act. When 
it ceases to act it can be restored to the active condition 
by passing an electric current through it in the opposite 
direction. Under these conditions the lead sulphate on 
one plate is reduced to metallic lead, and that on the other 
is oxidized to lead peroxide. These facts are taken 
advantage of in the preparation of the well-known storage 
batteries. 

Red Lead , Minium. — When lead oxide is gently heated 
in the air it takes up oxygen, and is converted into the red 
compound known as minium or red lead. The commer- 
cial article of this name varies in composition, but 
approximates to that represented by the formula Pb 3 4 , 
and, if the oxide is slowly heated, the amount of oxygen 
taken up is that required to form a compound of the above 
formula. Red lead varies in color from red to yellowish, 
according to the method of preparation. When heated 
it becomes dark, but the red color appears again on cool- 
ing. When heated to a high temperature it loses oxygen 
and yields lead oxide : 

Pb 3 4 = 3PbO + 0. 



588 COLLEGE CHEMISTRY. 

When treated with dilute nitric acid lead nitrate is 
formed, and*lead peroxide is left undissolved. Red lead 
is used as a pigment, and sometimes in place of litharge 
when an oxide of lead is needed, as in the manufacture 
of glass, as a flux in the manufacture of porcelain, etc. 

Lead Sulphide, PbS. — This has already been referred 
to as the principal compound from which lead is obtained. 
The natural variety is called galena or galenite. It is 
formed in the laboratory as a black precipitate, when 
hydrogen sulphide is passed into a solution of a lead salt. 
When heated in the air, as in the roasting of galenite, the 
sulphur passes off as sulphur dioxide, and the lead is con- 
verted into oxide. Concentrated hydrochloric acid dis- 
solves it. Concentrated nitric acid converts it into the 
sulphate. 

Lead Nitrate, Pb(N0 3 ) 2 . — The nitrate is easily made by 
dissolving lead, lead oxide, or carbonate in nitric acid. 
The salt crystallizes well, and is easily soluble in water. 
It is difficultly soluble in dilute nitric acid, and insoluble 
in concentrated nitric acid, resembling in this respect 
barium nitrate. It is decomposed by heat, giving nitrogen 
peroxide, N0 2 , and lead oxide. 

Lead Carbonate, PbC0 3 . — The carbonate occurs in 
nature as cerussite, crystallized in forms which are the 
same as those of barium carbonate, and of that variety of 
calcium carbonate known as aragonite. It can be obtained 
by adding lead nitrate to a solution of ammonium car- 
bonate, but, when solutions of lead salts are treated with 
the secondary carbonates of the alkali metals, precipitates 
of basic carbonates are always obtained. When an excess 
of sodium carbonate is added to a solution of lead nitrate, 
the precipitate has the composition 3Pb0.2C0 2 -f- H 2 0. 
The salts usually obtained are more complicated than this, 
but the relations between them and lead oxide and car- 
bonic acid are of the same kind. Basic lead carbonate is 
prepared and used extensively, under the name of white 
lead, as a pigment. It is manufactured by different 
methods. The principal ones are the following : 



LEAD CARBONATE- LEAD SULPHATE. 5 8 9 

(1) The Dutch Method. — This consists in exposing sheets 
of lead wound in spirals to the action of vinegar, air, and 
carbon dioxide from decaying organic matter. The spirals 
of sheet lead are placed in earthenware vessels, on the 
bottom of which, but not in contact with the lead, the 
vinegar is placed. The vessels thus arranged are placed 
in beds of horse-manure. In consequence of decomposi- 
tion, which is set up in the manure, carbon dioxide is 
given off slowly, and enough heat is generated to start the 
action upon the lead. The chemical changes involved in 
the process are, mainly, the formation of a basic acetate 
of lead, and the subsequent decomposition of this by 
carbon dioxide, forming a basic carbonate, and leaving 
the acetic acid free to act upon a further quantity of lead. 

(2) The French Method. — In this method a solution of 
basic lead acetate is prepared by treating a solution of the 
neutral salt with lead oxide. This is then decomposed by 
passing carbon dioxide into it, when a basic carbonate is 
thrown down. The carbon dioxide is generally made by 
burning coke. 

(3) The English Method. — This is a modification of the 
Dutch method, and differs from it chiefly in the substitu- 
tion of spent tan in a state of fermentation for manure, 
and the use of dilute acetic acid in place of vinegar. 
There is less risk of discoloration in consequence of the 
formation of hydrogen sulphide, but, the fermentation 
takes place more slowly, and the whole process, therefore, 
requires a longer time. 

An objection to white-lead paint is that it turns dark 
under the influence of hydrogen sulphide. It also turns 
yellow in consequence of the action of some substance 
contained in the oil with which the lead carbonate is 
mixed. 

Lead Sulphate, PbS0 4 , occurs to some extent in nature. 
It is formed by adding sulphuric acid or a soluble sulphate 
to a solution of a lead salt, and by oxidation of lead sul- 
phide. Like barium' sulphate, it is practically insoluble 
in water. As stated above, it is somewhat soluble in con- 



59° COLLEGE CHEMISTRY. 

centrated sulphuric acid, and it is therefore always found 
in the concentrated acid of commerce. Nitric acid and 
hydrochloric acid dissolve it in considerable quantity. It 
dissolves further quite readily in solutions of some am- 
monium salts, as in ammonium tartrate and acetate. 
When heated to redness it is partly decomposed with loss 
of sulphur trioxide. 

Reactions which are of Special Value in Chemical 
Analysis. — The reactions of lead salts with the soluble 
hydroxides, with sulphuric acid, hydrochloric acid, hy- 
drogen sulphide, soluble carbonates, potassium chromate 
and dichromate, are the ones which are principally used 
in analysis. All of these have been treated of in this 
chapter, with the exception of those with potassium 
chromate and bichromate, which will be taken up in the 
chapter on Chromium (which see). In anticipation it 
may be said that the reactions are based upon the fact 
that lead chromate, PbCr0 4 , like barium chromate, is in- 
soluble in water. 



The elements of Family V, Group A, are vanadium, 
columbium, didymium, and tantalum. As they are closely 
related to the members of Group B, of the same family, 
they were treated of at the end of Chapter XVIII in con- 
nection with the members of the phosphorus group. 
Among them the one which is least known is didymium. 
This in turn is more or less closely related to two other 
elements of nearly the same atomic weight which occur 
in Families III and TV. These are lanthanum and cerium. 
A few words in regard to these three rare elements will 
suffice for the present purpose. 

Lanthanum, Cerium, Didymium. 

These three elements occur together in several rare 
minerals of Norway, as cerite, gadolinite, and allanite. 

Cerite is a silicate of the three metals, and its composition. 



LANTHANUM, ETC. 59 T 

CO 



is represented by the formula La 4 V (Si0 4 ) 3 -f 3H 2 0. It 

is probably a mixture of three isomorphous silicates. The 
principal constituent is cerium silicate, Ce 4 (Si0 4 ) 3 . The 
perfect separation of the constituents of the mineral is a 
very difficult operation. 

Lanthanum, La (At. Wt. 138), forms an oxide of the 
formula La 2 3 , analogous to that of aluminium. Its 
chloride also is analogous to that of aluminium, and has 
the composition LaCl 3 ; and in all its salts it acts as a 
trivalent element. 

Cerium, Oe (At. Wt. 140), forms two series of com- 
pounds, in one of which it is trivalent, resembling lan- 
thanum, and the other members of the aluminium group; 
while in the other series it is quadrivalent, resembling 
silicon and the other members of the silicon group. 

Didymium, Di, has already been referred to on p. 339 
in connection with the members of Family V, Group A, 
which it resembles in some respects. In most of its com- 
pounds it is, however, trivalent, forming compounds, of 
some of which the following are the formulas : 

DiCl 3 , Di 2 3 , Di(N0 8 )„ Di 2 (S0 4 ) 8 , Di 2 (C0 3 ) 3 , etc 

Praseodymium, Pr, and Neodymium, Nd. — While the 
name didymium is still given above, and this substance 
dealt with as though it were an element, as it was at first 
held to be, it has been shown by Auer von Welsbach that 
it consists of two very similar elements to which he has 
given the names praseodymium and neodymium. When 
the double nitrate of ammonium and didymium is re- 
peatedly recrystallized it is separated into two salts, one 
of which is green, and the other rose-colored. When the 
nitrate or oxalate of one of these new elements is ignited 
it forms a black oxide, while from the other is formed an 
oxide of a different color. The element that gives green 
gaits is called praseodymium, and the other is called 



59 2 COLLEGE CHEMISTRY. 

neodymium. The atomic weight of neodymium is 143.6; 
that of praseodymium 140.5. 

EXPERIMENTS. 
Tin and its Compounds. 

Experiment 227.— Dissolve tin in hydrochloric acid and let the 
product crystallize. 

Experiment 228. — Pass dry chlorine over granulated tin con- 
tained in a retort connected with a receiver, using the arrange- 
ment illustrated in Fig. 60, page 318. Redistil the product. 
Treat some of the liquid with water, and boil. 

Lead and its Compounds. 

Experiment 229. -Make specimens of lead chloride and lead 
iodide, and crystallize them from water. 

Experiment 230. — Make lead sesquioxide by bringing together 
lead acetate and sodium hydroxide, and treating the solution 
with a solution of sodium hypochlorite. 

Experiment 231. — Treat some red lead with dilute nitric acid. 
Filter, wash, and treat the substance left on the filter with 
hydrochloric acid. 



CHAPTER XXXI. 

ELEMENTS OF FAMILY VI, GROUP A: 

CHROMIUM.— MOLYBDENUM.— TUNGSTEN.— 

URANIUM. 

General. — At the end of Chapter XIV, in connection 
with the elements of the sulphur group, the four elements 
which form the subject of this chapter were briefly referred 
to, for the reason that in some respects they resemble 
sulphur. As was there stated, this resemblance "is seen 
mainly in the formation of acids of the formulas H 2 Cr0 4 , 
H 2 Mo0 4 , H 2 W0 4 , and H 2 U0 4 ; and the oxides Cr0 3 , 
Mo0 3 , W0 3 , and U0 3 ." Further, it was stated that 
"when the acids of chromium, molybdenum, tungsten, 
and uranium lose oxygen, they form compounds that have 
little or no acid character. The lower oxides of chromium 
form salts with acids, and these bear a general resemblance 
to the salts of aluminium, iron, and manganese. The 
chromates lose their oxygen quite readily when acids are 
present with which the chromium can enter into combina- 
tion as a base-forming element." "Molybdenum and 
tungsten do not form salts of this character: indeed they 
seem to be practically devoid of the power to form bases. 
Uranium, on the other hand, forms some curious salts 
which differ from the simple metallic salts which we com- 
monly have to deal with. These are the uranyl salts which 
are regarded as acids, in which the hydrogen is either 
wholly or partly replaced by the complex U0 2 , which is 
bivalent. Thus the nitrate has the formula U0 2 (N0 8 ) 2 , 
the sulphate (U0 2 )S0 4 , etc. These salts are derived from 
the compound U0 2 (OH) 2 , acting as a base, whereas the 
compound has also distinctly acid properties." That 

593 



594 COLLEGE CHEMISTRY. 

member of the group the compounds of which are most 
commonly met with in the laboratory and in the arts is 
chromium, and this will receive principal attention here. 



Chkomium, Cr (At. Wt. 52.1). 

General. — This element forms three series of com- 
pounds, in which it appears to be respectively bivalent, 
trivalent, and sexivalent. Of these the members of the 
series in which it is trivalent are most stable under ordinary 
circumstances. Some of the principal members of the 
first series, or the chromous compounds, are represented by 
the formulas 



OrCl 2 , Cr(OH) 2 , CrS0 4 , CrC0 3 . 

Of the second series, or the chromic compounds, some of 
the principal members are : 

CrCl s , Cr 2 3 , Cr 2 (S0 4 ) 3 , Cr(N0 3 ) 3 , KCr(S0 4 ), + 12H 2 0. 

And, finally, the members of the third series are derived 
from the oxide Cr0 3 , and they are for the most part salts 
of the acid of the formula H 2 Cr0 4 , known as chromic acid, 
or of an acid of the formula H 2 0r 2 O 7 , known as Uchromic 
acid, which is closely related to chromic acid. 

When exposed to the air the chromous compounds are 
converted into chromic compounds, and they are in gen- 
eral readily converted into chromic compounds by the 
action of oxidizing agents, as cuprous and mercurous 
compounds are converted into cupric and mercuric com- 
pounds. If the oxidation takes place in acid solution the 
limit is reached when a chromic salt is formed. If, how- 
ever, the action takes place in the presence of a strong 
base the limit is reached in the formation of a chromate. 
Thus, suppose chromous oxide to be treated with an 
oxidizing agent in the presence of sulphuric acid, the final 



CHROMIUM. 595 

product would be chromic sulphate, as represented in the 
following equations : 

CrO + II 2 S0 4 = CrS0 4 + II 2 0; 

2CrS0 4 + H,S0 4 + = Cr 2 (S0 4 ) 3 + H 2 0. 

On the other hand, if the oxidation takes place in the 
presence of caustic potash the final product is potassium 
chromate, as shown in the following equation : 

CrO + 2KOII + 2 = K 2 Cr0 4 + H 2 0. 

When a chromate is treated with an acid it tends to pass 
back to a compound of the chromic series, and the change 
involves the giving up of oxygen. 

Forms in which Chromium occurs in Nature. — The 
principal form in which chromium occurs in nature is the 
mineral chromite, also known as chromic iron and chrome 
iron ore. This has the composition F'eCr 2 4 , and, as will 
be pointed out below, it is probably analogous to the 
spinels (see p. 526), being an iron salt of the acid CrO. OH, 
which may be called metachromous acid. This view is 

represented by the formula n r Ao>^ e - It occurs also 

in the mineral crocoisite, which is lead chromate, PbCr0 4 . 
The name chromium is derived from the Greek xP&>^ a > 
meaning color; and the element is so called because most 
of its compounds are colored. 

Preparation. — Chromium is prepared by heating chromic 
oxide with carbon in the electric furnace. The first 
product obtained is a mixture of carbides of chromium. 
When this is heated with lime in the electric furnace 
calcium carbide and metallic chromium are formed. By 
this method it is obtained in crystals. 

Properties. — Chromium has a bright metallic lustre and 
is capable of high polish. It is very hard and difficultly 
fusible. It burns brilliantly in oxygen, though it is not 
changed by exposure to the air at ordinary temperatures. 
It is dissolved by hydrochloric acid. Nitric acid does not 



59 6 COLLEGE CHEMISTRY. 

attack it. When treated with salts of potassium which 
easily give up their oxygen, as the chlorate and nitrate, it 
is converted into potassium chromate. 

Chromic Chloride, CrCl 3 . — This compound is made in 
solution by dissolving chromic hydoxide, Cr(OH) 3 , in 
hydrochloric acid. This solution has a dark-green color. 
When evaporated to a sufficient extent crystals of the com- 
position CrCl 3 -|- 6H 2 are deposited. If these are heated 
in the air they undergo decomposition just as aluminium 
chloride does, and the product left behind is chromic 
oxide : 

2CrCl 3 + 3H 2 = Cr 2 3 + 6HC1. 

If, however, the crystallized chloride is heated in an 
atmosphere of chlorine or hydrochloric acid, the water is 
given off, and the anhydrous chloride, which has a beauti- 
ful reddish-violet color, is formed. This dissolves in water 
and forms a green solution. But if the dry chloride thus 
obtained is sublimed, it is deposited in lustrous laminae of 
the same color; and this variety is insoluble in water and 
acids, and is only slowly acted upon by boiling alkalies. 
This insoluble, crystallized variety of the chloride is 
obtained also by the same method as that used in making 
aluminium chloride, that is, by passing a current of chlo- 
rine over a heated mixture of carbon and chromic oxide. 

Chromic Hydroxide, Cr(OH) 3 . — When ammonia is 
added to a solution of a chromic salt, a light-blue volu- 
minous precipitate, which has the composition Cr(OH) 3 -}- 
2H 2 0, is formed. When this is filtered off and dried in a 
vacuum it loses the water and leaves the hydroxide. This 
is readily converted by heat into a compound of the 
formula CrO. OH, and finally into chromic oxide, Cr 2 3 . 
The green precipitates formed in solutions of chromic 
salts by sodium and potassium hydroxides always contain 
some of the alkali metal in combination. 

Chromic hydroxide, like aluminium hydroxide, dissolves 
in the soluble hydroxides, and forms salts known as 
chromites, which are derived from the acid CrO. OH, 



CHROMIC OX IDF. -CHROMIC SULPHATE. 597 

Thus with potassium hydroxide the action takes place as 
represented in the equation 

/OB /n 

Or^-OH + KOH = Cr^ K + 2II 2 0. 
\OH UJ ^ 

When a solution containing potassium or sodium ehromite 
is boiled the salt is decomposed and chromic hydroxide 
precipitated, though the precipitate thus formed always 
(vi 11 tains some of the alkali metal in combination. It will 
be noticed that in this respect aluminium and chromium 
conduct themselves differently towards the alkaline hy- 
droxides. 

It has already been stated that ehromite, (CrO.O) 2 Fe, 
is regarded as an iron salt of the same order as the potas- 
sium salt referred to. 

Chromic Oxide, Cr 2 3 , is formed by igniting the hy- 
droxides, and is most readily prepared by heating a mixture 
of potassium dichromate and sulphur. The sulphur is 
oxidized, and with the potassium forms potassium sul- 
phate, while the chromic acid is reduced to the form of 
the oxide Cr 2 3 . It can be obtained in crystals. As 
ordinarily obtained it is a green powder, which after igni- 
tion is almost insoluble in acids. It is, however, dissolved 
by treatment with fusing acid potassium sulphate. The 
oxide colors glass green, and is used in painting porcelain. 

Chromic Sulphate, Cr 2 (S0 4 ) 3 , is made by dissolving the 
hydroxide in concentrated sulphuric acid when it is 
deposited in purple crystals of the composition Cr 2 (S0 4 ) 3 -f- 
15H 2 0. If the solution of this salt is boiled, the solution 
becomes green, and crystals cannot be obtained from it. 
But by standing for some time the green solution becomes 
reddish purple again, and yields the crystallized salt. 
Other salts of chromium act in the same way. They exist 
in two varieties, one of which crystallizes and is reddish- 
purple in color, while the other does not crystallize and is 
green. The crystallized salts are converted into the un- 
crystallized green salts by boiling, and the green salts are 
converted into the crystallized salts by standing. 



■■ 



59 § COLLEGE CHEMISTRY. 

Chrome-Alums. — Chromic sulphate, like aluminium 
sulphate, combines with other sulphates, such as potas- 
sium, sodium, and ammonium sulphates, and forms well- 
crystallized salts, which are closely analogous to ordinary 
alum. They all contain twelve molecules of water, as 
represented in the formulas below: 

Chrome- Alum KCr(S0 4 ) 2 + 12H 2 

Sodium Chrome- Alum NaCr(S0 4 ) 2 + 12H 2 

Ammonium Chrome-Alum (NHJCr(S0 4 ) 2 + 12H 2 

The potassium compound which is commonly called 
chrome-alum is made by adding a reducing agent, such as 
alcohol or sulphur dioxide, to a solution of potassium 
dichromate containing sulphuric acid. If the solution is 
heated it turns green, and crystals cannot be obtained 
from it. But on standing for a considerable time its color 
changes, and reddish-purple crystals of the alum are 
deposited. This change can be facilitated by putting 
some crystals of the salt in the concentrated green solu- 
tion. The action of reducing agents upon potassium 
dichromate will be treated of farther on. The salt finds 
application in dyeing and tanning. 

Chromic Acid and the Chromates. — It has been stated 
that when chromium compounds belonging to the chro- 
mous and chromic series are oxidized in the presence of 
bases they are converted into chromates. These salts 
are derived from an acid of the formula H 2 Cr0 4 , which is 
unknown, as it breaks down spontaneously into chromium 
trioxide, Cr0 3 , and water, when it is set free from its 
salts, just as carbonic and sulphurous acids break down 
respectively into carbon dioxide and water, and sulphur 
dioxide and water. The starting-point for the preparation 
of the chromates and the compounds related to them is 
chromic iron. This is ground fine, intimately mixed with 
a mixture of caustic potash and lime, and then heated in 
shallow furnaces in contact with the air. Under these 
circumstances oxidation is effected by the oxygen of the 
air. The iron is converted into ferric oxide, and the 



POTASSIUM CH ROM ATE, ETC. 599 

chromium gives, with the calcium and potassium, the 
corresponding chromates, CaCr0 4 and K 2 Cr0 4 . When the 
mass is treated with water these salts dissolve, and ferric 
oxide remains undissolved. By treating the solution with 
potassium sulphate the calcium salt is converted into the 
potassium salt, and thus all the chromium appears in the 
form of potassium chromate. The changes referred to 
are represented in the following equations : 

2(Cr0 2 ) 2 Fe + 8KOH -f 70 = 4K 2 Cr0 4 + Fe 2 3 + 4H 2 0; 
2(Cr0 2 ) 2 Fe + 4CaO + 70 = 4CaCr0 4 + Fe 2 3 ; 
CaCr0 4 -|-K 2 S0 4 = K 2 Cr0 4 + CaS0 4 . 

As potassium chromate is easily soluble in water, and 
therefore difficult to purify, it is converted into the bichro- 
mate, which is less soluble and crystallizes well. The 
change is easily effected by adding the necessary quantity 
of a dilute acid. If nitric acid is used the reaction is 
represented by the following equation : 

2K 2 Cr0 4 + 2HN0 3 = K 2 0r 2 7 + 2KN0 3 + H 2 0. 

The salt thus obtained is manufactured on the large scale 
and is the starting-point for the preparation of other 
chromium compounds. 

Potassium Chromate, K 2 Cr0 4 , formed as above 
described, is a light-yellow crystallized substance which is 
easily soluble in water. It is isomorphous with potassium 
sulphate. Acids convert it into the bichromate, as just 
stated. 

Potassium Bichromate, K 2 Cr 2 7 . — This salt forms large 
triclinic red plates. It is soluble in ten parts of water 
at the ordinary temperature, and is much more soluble 
in hot water. When heated it at first melts without 
undergoing decomposition; at white heat, however, it is 
decomposed, yielding the chromate, chromic oxide, and 
oxygen : 

2K 2 Cr 2 7 = 2K 2 Cr0 4 + Cr 2 3 + 30. 



6oo COLLEGE CHEMISTRY. 

It undergoes a similar change, but much more readily, 
when heated with concentrated sulphuric acid. In this 
case, however, the chromic oxide forms chromic sulphate 
with the acid, and this forms chrome-alum with the potas- 
sium sulphate : 

K 2 Cr 2 7 + 4H 2 S0 4 = 2KCr(SOJ 2 + 4H 2 + 30. 

All the oxygen in the chromate in excess of that required 
to form the alum and water is given off. This also is the 
character of the action towards reducing agents in general. 
Concentrated hydrochloric acid is oxidized by the 
dichromate, and chlorine is evolved: 

K 2 Cr 2 7 + 14HC1 = 2KC1 + 2CrCl 3 + 7H 2 + 601. 

Here two atoms of chlorine are required to form potassium 
chloride with the potassium, and six to form chromic 
chloride with the chromium; and the eight hydrogen 
atoms in combination with this chlorine combine with four 
atoms of oxygen of the dichromate, leaving three more to 
oxidize hydrochloric acid. Consequently one molecule of 
the dichromate sets free six atoms of chlorine : 

30 + 6H01 = 3H 2 + 601. 

When the bichromate in solution is treated with potas- 
sium hydroxide, its color changes to yellow, in consequence 
of the formation of the chromate, the action taking place 
as represented in this equation : 

K 2 Cr 2 7 + 2K0H = 2K 2 Cr0 4 + H a 0. 

Potassium bichromate finds extensive use in the arts 
and in the laboratory as an oxidizing agent. With gelatin 
it forms a mixture sensitive to light, which turns it dark, 
and makes it insoluble. This fact is made the basis of 
a number of photogravure or photographic reproduction 
processes. The bichromate is used, further, in dyeing. 

Chromium Trioxide, Cr0 3 , crystallizes out on cooling 
when either the chromate or the bichromate is treated in 



CHROMATES AND BICHROMATES. 60 1 

concentrated solution with concentrated sulphuric acid. 
It is a beautiful red substance, which crystallizes in needles. 
When dissolved in water it forms a solution from which, 
by neutralization, the chromates can be obtained. It is 
an extremely active oxidizing agent, disintegrating most 
organic substances with which it is brought in contact. 

Relations between the Chromates and Bichromates. — 
The fact that chromium trioxide with water gives chromic 
acid, which is a dibasic acid, whose salts in general resem- 
ble those of sulphuric acid, leads to the belief that the 
structure of chromic acid should be represented by a 
formula similar to that of sulphuric acid, thus : 

2 S(OH), O.Cr(OH), 

Just as sulphuric acid by loss of water is converted into 
disulphuric acid or pyrosulphuric acid, so chromic acid is 
converted into bichromic acid, and in all probability the 
relation between the chromates and bichromates is the 
same as that between the sulphates and disulphates, as 
represented by the equations 

0H .OH 



SO *<OH S0 3 <^ H 
C ^<oi_Cr0 2 <OH 



But, as has been stated, neither chromic acid nor bi- 
chromic acid is known, as they break down into chromium 
trioxide and water when set free from their salts. The 
conversion of potassium chromate into the bichromate by 
treatment with an acid is represented as shown below: 

. A OK n A ^OK 

Cr ° 2 <OK,HN0 3 Cr ° 2< OH , 2KN0 . 
r n ^OK + HNO = Prn OH + 2KN °3> 
Cr0 2 < OK 3 Cl ° 2 <OK 



602 COLLEGE CHEMISTRY. 

Sodium Chromate, Na 2 Cr0 4 , and Sodium Bichromate, 

Na 2 Cr 2 7 , are made in the same way as the potassium 
compounds. They are both deliquescent as ordinarily 
made. It has, however, recently been shown that a 
sodium bichromate which is not deliquescent can be made; 
and at present it is largely manufactured instead of the 
somewhat more expensive potassium salt. 

Barium Chromate, Ba0r0 4 , like the sulphate, is insolu- 
ble in water, and is precipitated when a solution of a 
barium salt is brought together with a soluble chromate 
or bichromate. It is soluble in hydrochloric acid and 
nitric acid, but not in acetic acid. The strontium salt is 
somewhat soluble in water, and easily in hydrochloric, 
nitric, and acetic acids. 

Lead Chromate, Pb0r0 4 , occurs in nature, and is 
formed as a beautiful yellow precipitate when a solution 
of a lead salt is treated with a solution of a chromate or 
bichromate. It is used as a pigment under the name 
chrome yellow. 

Reactions which are of Special Value in Chemical 
Analysis. — The reactions of chromic salts with the alka- 
line hydroxides have been explained. With the soluble 
carbonates they give precipitates which consist mainly of 
the hydroxide, though carbonic acid is to some extent in 
combination in them. These precipitates are soluble in a 
large excess of the carbonate. 

The solution of chromic oxide in an alkali is green, but 
by oxidizing agents it is turned yellow in consequence of 
the formation of a chromate. 

Hydrogen sulphide does not precipitate chromium from 
its salts. Ammonium sulphide precipitates chromium 
hydroxide, the reaction being the same as in the case of 
aluminium (which see). 



MOLYBDENUM. 603 

Chromates give with barium and lead salts yellow pre- 
cipitates (see Barium Chromate and Lead Chr ornate). 

When heated with sulphuric acid, the chromates are 
decomposed, with evolution of oxygen. The action of 
hydrochloric acid upon the chromates was explained above. 
With borax or microcosmic salt chromium compounds 
form green beads, both in the oxidizing and reducing 
flames. 



Molybdenum, Mo (At. Wt. 96). 

General. — Molybdenum is of interest on account of the 
variety of its compounds. It forms four compounds with 
chlorine, the formulas of which appear to be MoCl 2 , 
M0CI3, MoCl 4 , and MoCl 5 . With oxygen also it forms 
four compounds, but, while the first three are analogous to 
the first three chlorine compounds above mentioned, the 
last differs from the last chlorine compound. In it the 
element is sexivalent. The formulas of the oxygen com- 
pounds are MoO, Mo 2 3 , Mo0 2 , and Mo0 3 . The last of 
these is analogous to chromium trioxide, as it forms salts 
with bases analogous to the chromates, and known as the 
molybdates. 

Occurrence and Preparation. — Molybdenum occurs in 
nature principally as molybdenite, which is the sulphide 
MoS 2 , and as wulfenite, which is lead molybdate, PbMo0 4 . 
It occurs also, but in smaller quantity, as molybdenum 
trioxide, Mo0 3 . It is obtained in free condition by heat- 
ing the oxides or chlorides in a current of hydrogen. 

Molybdic Acid and the Molybdates.— Molybdenum tri- 
oxide, Mo0 3 , combines readily with bases, forming salts 
which in composition are analogous to the chromates. 
When ammonium molybdate is treated with moderately 
dilute nitric acid free molybdic acid crystallizes out. 
This has the composition represented by the formula 
H 2 Mo0 4 + H 2 0. 

Lead Molybdate, PbMo0 4 , occurs in nature, as has been 
stated, and the mineral is known as wulfenite. It can be 



604 COLLEGE CHEMISTRY. 

obtained artificially by melting together sodium molybdate, 
lead chloride, and sodium chloride; or by treating a solu- 
tion of sodium molybdate with a solution of lead nitrate. 
If the reagents are pure the artificially prepared salt is 
white, while the natural variety is always yellow or red. 

Phospho-molybdic Acid. — Among the best known and 
most frequently met with compounds of molybdic acid 
with acids is that which it forms with phosphoric acid, 
known as phospho-molybdic acid. When a solution of 
ammonium molybdate in an excess of nitric acid is added 
in excess to a solution of phosphoric acid or a phosphate, 
a yellow precipitate is formed. This is ammonium phospho- 
molybdate, which, when dried, has the composition repre- 
sented by the formula 12Mo0 3 .(NH 4 ) 3 P0 4 . This is 
insoluble in water and in dilute acids, and also in a nitric- 
acid solution of ammonium molybdate. On account of 
the properties mentioned, this salt furnishes a valuable 
means of detecting phosphoric acid and of precipitating it 
from its solutions. When the salt is treated with aqua 
regia it is decomposed, and from the solution formed a 
compound of the composition H 3 P0 4 .llMo0 3 + 12H 2 
crystallizes out. 

Tungsten, W (At. Wt. 184). 

General. — Like molybdenum, tungsten forms a large 
variety of compounds. With chlorine it forms four, of 
which the formulas are WC1 2 , WC1 4 , WC1 5 , and W01 6 . 
With oxygen, however, it forms but two compounds, and 
these are represented by the formulas W0 2 and W0 3 . The 
trioxide forms salts with bases which are analogous to the 
molybdates. 

Occurrence and Preparation. — Tungsten occurs in 
nature as tungstates. The principal one is the iron salt, 
which always, however, contains some manganese. This 
is known as wolframite, and has the composition repre- 
sented by the formula FeW0 4 . Calcium tungstate, 
CaW0 4 , or scheelite, and lead tungstate, PbW0 4 , or 



TUNGSTEN, URANIUM. 605 

stolzite, are also found in nature, but in smaller quantity 
than wolframite. 

Properties. — Tungsten forms lustrous, steel-colored 
laminae, or a black powder. It is very hard and difficultly 
fusible, and lias the specific gravity 19.129. It is not 
changed by contact with the air at ordinary temperatures. 
At higher temperatures it combines with oxygen and forms 
the trioxide, AV0 3 . Nitric acid and aqua regia convert it 
into the trioxide. It is used in the manufacture of steel, 
as the addition of from 8 to 9 per cent of it makes steel 
extremely hard. 

Tungstic Acid and the Tungstates. — When the required 
quantities of tungsten trioxide and potassium carbonate 
are brought together in solution, or are melted together, 
potassium tungstate, K 2 W0 4 , is formed. If a solution of 
this salt is treated with a strong acid at the ordinary tem- 
perature, a white precipitate of the composition H 2 W0 4 -(- 
H 2 is formed. This is tungstic acid, analogous to 
crystallized molybdic acid. If the solutions are hot the 
precipitate has the composition H 2 W0 4 . 

Uranium, U (At, Wt. 239.5). 

General. — Uranium has stronger basic properties than 
either molybdenum or tungsten; and it differs from 
chromium in the fact that the trioxide forms salts with 
acids. These salts are the uranyl salts which are derived 
from the hydroxide, U0 2 (OH) 2 or H 2 U0 4 . The corre- 
sponding compounds of chromium, molybdenum, and 
tungsten are acids. Uranium also forms salts in which it 
acts as a quadrivalent element, as U(80 4 ) 2 . While the 
hydroxide, U0 2 (OH) 2 , forms salts with acids, it also forms 
salts with the strongest bases. These are analogous in 
composition to the bichromates, and have the general 
formula M 2 U 2 7 . With chlorine, uranium forms the 
compounds UC1 3 , UC1 4 , and UCL; and with oxygen the 
following: U0 2 , U 3 8 , U0 3 , and U0 4 . 

Occurrence and Preparation. — Uranium occurs in nature 
chiefly in the form of the mineral known as pitch-blende 



606 COLLEGE CHEMISTRY. 

or uraninite, which consists of the oxide, U 3 8 , mixed 
with a number of other substances. When this is finely 
powdered and treated with concentrated nitric acid, uranyl 
nitrate, U0 2 (N0 3 ) 2 , is obtained, and, by igniting this at 
not too high a temperature, the trioxide, U0 3 , is left 
behind. In order to isolate the metal, the oxide, U 3 8 , 
is heated with charcoal in the electric furnace. 

Properties. — Uranium has the color of nickel and the 
specific gravity 18.4. 

Oxides. — The oxide of uranium which is formed as the 
last product of the action of oxygen on uranium or the 
other oxides when these are heated in the air is that which 
has the composition U 3 8 , which is also the composition 
of the natural variety. When this is treated with nitric 
acid, however, it is converted into uranyl nitrate, 
U0 2 (N0 3 ) 2 , which is a derivative of the trioxide, U0 3 ; 
and when the nitrate is ignited, the trioxide is left behind. 
By reduction with hydrogen the trioxide is converted into 
the dioxide, U0 2 ; and when either the dioxide or the tri- 
oxide is heated in the air, the product obtained is the 
oxide U 3 8 . 

Uranous Salts. — In the uranous salts, uranium acts as 
a quadrivalent element, replacing four atoms of hydrogen, 
as, for example, in the sulphate, which has the composi- 
tion U(S0 4 ) 2 . But few salts of this order are known. 

Uranyl Salts. — As already explained, the uranyl salts 

OTT 

are derivatives of the hydroxide U0 2 <^tt, and are formed 

by the action of acids, as represented in the equations 
below : 

^^OH + HO >SO * = U ° 2 <0 >S ° 2 + 2H *°; 

J0 OH HO.NO, _ UQ O.NO, H 
u ° 2 < OH ^ HO. NO, ~ U U * < 0. NO, + Zt± » U ' 

They are derived from the acids by substituting uranyl, 
U0 2 , which is bivalent, for the hydrogen. — Uranyl 
nitrate, U0 2 (N0 3 ) 2 , is easily obtained, as above described. 



URA NATES. 607 

and crystallizes well in lemon-yellow prisms. — Uranyl 
sulphate, U0 2 (S0 4 ), is formed by treating the nitrate with 
sulphuric acid. 

Uranates. — When a uranyl salt is treated with a soluble 
hydroxide a precipitate is formed which is a salt of an 
acid, H 2 U 2 7 , which maybe called diuranic acid, as in 
composition it is analogous to bichromic and disulphuric 
acids. 

Sodium diuranate, Na 2 U 2 7 , is a fine yellow powder, 
which is manufactured and sold under the name uranium 
yelloiv, being used as a pigment for coloring glass, etc. — 
Ammonium diuranate, (NH 4 ) 2 U 2 7 , is also manufactured 
on the large scale. When it is treated with a solution of 
ammonium carbonate it dissolves, and from the solution 
a salt of the composition U0 2 (C0 3 ) + 2(NHJ 2 C0 3 crystal- 
lizes out. The solubility of ammonium diuranate in am- 
monium carbonate is utilized in analysis. 

Many uranium salts exhibit in solution a beautiful 
fluorescence. 



EXPERIMENTS. 
Chromic Acid and the Chromates. 

Experiment 232.— Powder some chromic iron very fine. Add 
3 grams to a molten mixture of 3 grams each of potassium car- 
bonate, potassium hydroxide, and potassium nitrate, heated in 
a porcelain crucible. After cooling treat the mass with water. 
Potassium cbromate is in the solution. 

Experiment 233. — To the solution of potassium cbromate ob- 
tained in the last experiment add nitric acid to decompose the 
unacted-upon potassium carbonate, and give the solution an 
acid reaction. The color will change from yellow to red. The 
red color indicates the presence of the bichromate. 

Experiment 234. — Treat a solution of 10 to 20 grams potas- 
sium bichromate with potassium hydroxide until the color be- 
comes pure yellow, and evaporate to crystallization. 

Experiment 235. — Make a solution of potassium bichromate 
saturated at the ordinary temperature. Pour into this 1| times 
its volume of ordinary concentrated sulphuric acid. After the 



608 COLLEGE CHEMISTRY. 

liquid cools, and the chromium trioxide separates, filter with 
the aid of a filter-pump through glass-wool. 

Experiment 236. — To a solution of potassium bichromate add 
some hydrochloric acid and a little alcohol. On boiling, the 
alcohol is oxidized, and the solution now contains chromic 
chloride. 



CHAPTER XXXII. 

ELEMENTS OF FAMILY VII, GROUP At 
MANGANESE (Mn, At. Wt. 55). 

General. — At the close of Chapter XII (which see), 
which treated of the elements of Family VII, Group B, 
or the chlorine group, reference was made to manganese, 
and attention was called to the fact that in some respects 
it resembles chlorine. The resemblance is seen in the 
formation of an oxide, Mn 2 0_ , and an acid, HMn0 4 , 
analogous to perchloric acid. On the other hand, in most 
of its compounds it plays the part of a base-forming 
element, and in this capacity it forms two series of com- 
pounds, known as the manganous and the manganic com- 
pounds. In the former the element appears to be bivalent, 
and in the latter trivalent. The formulas of some of the 
principal manganous compounds are: 

MnCl 2 , Mn(OH) 2 , MnO, Mn(NO,) a , MnS0 4 , MnC0 3 , etc. 

The formulas of some of the principal manganic com- 
pounds are: 

Mn(OH) 3 ,Mn 2 3 ,MnO(OH),Mn 2 (S0 4 ) 3 , 
KMn(SOj 2 + 12H 2 0, etc. 

These two series of compounds are analogous in composi- 
tion to the chromous and chromic compounds, but, while 
the chromic compounds are more stable than the chromous 
compounds, the manganous compounds are more stable 
than the manganic compounds. B}^ contact with the air, 
the manganous are * not as a rule converted into the man- 
ganic compounds. 

609 



6 to COLLEGE CHEMISTRY. 

Corresponding to chromic acid, there is a manganic 
acid, H 2 Mn0 4 ; and, further, there is the permanganic acid 
already mentioned, of the formula HMn0 4 . An analogous 
compound of chromium, perchromic acid, HCr0 4 , is 
believed to be formed when hydrogen peroxide is added 
to an aqueous solution of chromic acid, but this is by no 
means certain. Manganic acid and its salts are very un- 
stable, and are readily converted into permanganic acid 
and the permanganates. On the other hand, perchromic 
acid, if it exists at all, is spontaneously decomposed, yield- 
ing ordinary chromic acid. To sum up, then, both 
chromium and manganese form four classes of compounds. 
But, while chromium under ordinary circumstances prefer- 
ably forms chromic salts and salts of chromic acid, man- 
ganese preferably forms manganous salts and salts of 
permanganic acid. 

Manganese forms a number of oxides corresponding to 
the formulas MnO, Mn 3 3 , Mn 3 4 , Mn0 2 . and Mn 2 0_. 

Forms in which Manganese occurs in Nature. — The 
principal natural compound of manganese is the black 
oxide or pyrolusite, MnO,. Besides this, however, there 
are several manganese compounds found in nature, the 
principal ones being braunite, Mn 2 3 , hausmannite, 
Mn 3 4 , manganite, Mn 2 2 (OH) 2 , and rhodocroisite, which 
is the carbonate, MnC0 3 . 

Preparation and Properties. — The metal is isolated from 
its oxides by heating them to a high temperature with 
charcoal in the electric furnace. It looks like cast iron, 
is brittle and hard, and has the specific gravity 8. It 
easily becomes oxidized in the air, decomposes warm 
water, and dissolves readily in dilute acids. It is used as 
a constituent of some useful alloys, and imparts certain 
desirable properties to iron, as will be pointed out when 
that metal is taken up. 

Manganous Chloride, MnCl.,. — This chloride is obtained 
in solution by dissolving any one of the oxides or hy- 
droxides or the carbonate of manganese in hydrochloric 
acid with the aid of gentle heat. Its formation in the 



OXIDES OF MANGANESE. 611 

preparation of chlorine from manganese dioxide and 
hydrochloric acid was referred to under Chlorine (which 
see). 

When manganic hydroxide, Mn(OH) 8 , is treated in the 
cold with hydrochloric acid, a deep brown-colored solution 
is formed, which is believed to contain the trichloride, 
M11CI3. On standing, however, this solution gives off 
chlorine slowly, and when heated it gives it off rapidly, 
and the color changes to pink when only manganous 
chloride is left in the solution. 

Phenomena similar to those just mentioned are observed 
when manganese dioxide is treated with hydrochloric acid 
in the cold, and it is believed that the tetrachloride, MnCl 4 , 
is contained in the solution. 

While the tetrafluoride itself has not been isolated, a 
solution is obtained by treating manganese dioxide with 
concentrated hydrofluoric acid which with potassium 
fluoride gives a salt of the formula MnF 4 .2KF or K.,MnF 6 . 

General Remarks concerning the Oxides. — The series 
of oxides of manganese strongly suggests that of the oxides 
of lead. Manganese, however, forms one oxide, the 
heptoxide, Mn,0. , of which there is no analogue among 
the compounds of lead. Placing the formulas of the oxides 
of the two metals side by side, we have the following table : 
PbO MnO 

Pb 2 3 Mn a 3 

Pb,0 4 Mn 3 4 

PbO., MnO., 

Mn/)_. 
Just as lead oxide when heated is converted into red-lead, 
Pb 3 4 , so the other oxides of manganese are converted into 
the oxide, Mn 3 4 , when heated. This has already been 
seen in the preparation of oxygen by heating the dioxide. 
In the same wa}~, the oxide, Mn 3 O s , loses enough oxygen, 
and the lowest oxide, MnO, takes up enough to form the 
same product: 

3Mn„0 3 = :>Mn„0 4 -f 0; 
3MnO + = Mn 3 4 . 



6 I2 COLLEGE CHEMISTRY. 

When treated with energetic oxidizing agents in the 
presence of the alkalies, all the oxides are converted into 
manganates. On the other hand, if a manganate is 
reduced in the presence of an acid the tendency to form 
manganous compounds shows itself, and all oxygen present 
in excess of that required to form the manganous salt is 
given off. 

Manganous Hydroxide, Mn(OH) 2 , is formed as a white 
precipitate when a soluble hydroxide is added to a solution 
of a manganous salt. Suspended in the alkali, or in con- 
tact with air, it absorbs oxygen, and is converted into 
hydroxides corresponding to the higher oxides. 

Manganous-manganic Oxide, Mn 3 4 , occurs in nature 
as the mineral hausmannite. It is formed, as already 
stated, by igniting the other oxides in contact with the 
air. When heated with dilute nitric acid it acts like the 
corresponding oxide of lead, giving manganous nitrate, 
and leaving manganese dioxide : 

Mn 3 4 + 4KN0 3 = 2Mn(N0 3 ) 2 + Mn0 2 + 2H 2 0. 

It breaks down in the same way with dilute sulphuric acid. 

Manganic Oxide, Mn 2 3 , occurs in nature as the min- 
eral braunite, and it can be made from the other oxides 
by igniting them in oxygen. A hydroxide related to this, 
and having the composition MnO.OH, analogous to the 
compounds of aluminium and chromium of the formulas 
AIO.OH and CrO.OH, is found in nature, and is known 
as manganite. The hydroxide, Mn(OH) 3 , is formed when 
manganous hydroxide, Mn(OH) 2 , is exposed in a solution 
of ammonia in contact with the air, and forms a brownish- 
black powder. 

Manganese Dioxide, Mn0 2 . — This important compound 
occurs in nature in considerable quantity, and is known 
as pyrolusite or the black oxide of manganese. It is 
obtained artificially by gently igniting manganous nitrate. 
A hydroxide derived from the dioxide is obtained by treat- 
ing a manganous salt in alkaline solution with a soluble 



WELDON'S PROCESS. 613 

hypochlorite or chlorine or bromine. The chiei applica- 
tion of the dioxide is in the preparation of chlorine. It 
is also used for making oxygen and for the purpose of 
decolorizing glass. In the last process a small quantity 
is added to the molten glass. This alone would give the 
glass an amethyst color. Without it the glass would be 
green. One color counteracts the other, and the glass 
appears colorless. 

Weldon's Process for the Regeneration of Manganese 
Dioxide in the Preparation of Chlorine. — Under the head 
of Chlorine (which see), Weldon's process was referred to; 
but as a satisfactory explanation of the working of the 
process could not be given without dealing with some 
rather complicated compounds of manganese, a fuller 
account was postponed until these compounds should be 
taken up. The object in view is to utilize the waste 
liquors from the chlorine factories. When manganese 
dioxide is treated with hydrochloric acid, as we have seen, 
manganous chloride and chlorine are formed, according 
to the equation 

Mn0 2 + 4HC1 = MnCl 2 + Cl 2 + 2H 2 0. 

The manganous chloride was of no special value until it 
was shown that by a comparatively simple method it can 
be converted into 'a compound which with hydrochloric 
acid gives chlorine. When it is treated in solution with 
lime the corresponding hydroxide is precipitated : 

MnCl 2 + Ca(OH) 2 = Mn(OH) 2 -f CaCl 2 ; 

and when this hydroxide mixed with lime is allowed to 
stand exposed to the air oxidation takes place, and a com- 
pound CaMn0 3 or CaMn 2 5 is formed : 

Mn(OH), + Ca(OH) 2 + = CaMn0 3 + 2H 2 0; 
2Mn(OH) 2 + Ca(OH), + 20 = CaMn 2 5 + 3H 2 0. 



6 14 COLLEGE CHEMISTRY. 

These compounds give chlorine when treated with hydro- 
chloric acid. They may indeed be regarded as consist- 
ing of lime and manganese dioxide, CaO.Mn0 2 and 
Ca0.2Mn0 2 , and the action of hydrochloric acid takes 
place thus : 

CaO.Mn0 2 + 6HC1 = CaCl 2 + MnCl 2 + 3H 2 + Cl 2 ; 
Ca0.2Mn0 2 + 10HC1 = CaCl 2 + 2MnCl 2 -f 5H 2 + 201 2 . 

In practice the waste liquor is mixed with calcium car- 
bonate in order to neutralize the acid. After settling, 
lime enough is added to precipitate the manganese as 
hydroxide, and to form with this a mixture in molecular 
proportions. Into this mixture steam and air are passed, 
when the oxidation referred to takes place, and calcium 
manganite is formed. 

Sulphides. — When a solution of a manganous salt is 
treated with ammonium sulphide, a flesh-colored precipi- 
tate, which is thought to be the hydrosulphide, is formed. 
When this is exposed to the air it turns dark in conse- 
quence of oxidation; and if allowed to stand in the liquid, 
if this is concentrated, it turns green and becomes crystal- 
line. The product thus formed is manganous sulphide, 
MnS. This occurs in nature as alabandite. A disulphide, 
MnS 2 , corresponding to the dioxide is also found in 
nature, and is known as hauerite. 

Manganous Sulphate, MnS0 4 , is formed by heating the 
oxides of manganese with concentrated sulphuric acid. 
If higher oxides than manganous oxide are used, oxygen 
is given off: 

MnO +H 2 S0 4 =MnS0 4 +H 2 0; 
Mn 2 3 + 2H 2 S0 4 = 2MnS0 4 + 2H 2 + 0; 
Mn0 2 + H 2 S0 4 = MnS0 4 + H 2 + 0. 

It crystallizes at low temperatures with seven, molecules of 
water, and at ordinary temperatures with five, in this 
respect resembling cupric sulphate (which see). The salt 
of the formula MnS0 4 -f 7H f orms bright-red monoclinic 



MANGANIC SULPHA TE t ETC. 6r5 

prisms; while that of the formula MnS0 4 + 5H a Conns 
pink triclinic crystals. Between 20° and 30° it forms 
monoclinic prisms with four molecules of water. 

Manganic Sulphate, Mn a (S0 4 ) 8 , is formed when the 
oxide Mn 3 4 , or the finely divided precipitated dioxide, 
MnO., , is treated with sulphuric acid at not too high a 
temperature. It forms a dark-green amorphous powder, 
which is easily decomposed by heat and by water. With 
the sulphates of the alkali metals it forms salts analogous 
to the alums, as KMn(S0 4 ) a + 12H 2 and NH 4 Mn(S0 4 ) a 
+ 12H„0, in which manganese takes the place of alumin- 
ium. 

Manganic Acid and the Manganates. — When an oxide 
of manganese is treated with an energetic oxidizing agent 
in the presence of a strong base it is converted into a 
manganate, just as the oxides of chromium are converted 
into chromates, and the compounds of sulphur into sul- 
phates. These three classes of compounds are analogous 
as far as the composition is concerned, as shown by the 
formulas 

M 2 Mn0 4 M 2 Cr0 4 M 2 S0 4 

Manganate Chromate Sulphate 

The manganates are, however, quite unstable except in 
alkaline solution, and when they decompose they form 
the permanganates. — Potassium manganate, K 2 Mn0 4 , is 
formed by fusing manganese dioxide with potassium 
hydroxide, when, if the air is not in contact with the mass, 
the reaction takes place as represented in the equation 

3Mn0 2 + 2KOH = K 2 Mn0 4 + Mn a O, + H 2 0. 

It is also made by .fusing the dioxide with potassium 
hydroxide and potassium chlorate, when this reaction takes 
place : 

3Mn0 2 + (5KOH 4- KC10 3 = 3K 2 Mn0 4 -f KC1 + 3H 2 0. 

When the mass obtained in either way is treated with 
water a dark-green solution of the manganate is formed, 



616 COLLEGE CHEMISTRY. 

and by allowing this to evaporate at the ordinary tempera- 
ture in a partial vacuum, or in an atmosphere free of 
oxygen, the salt is obtained in small crystals, which are 
almost black. When a solution of a manganate is treated 
with an acid, the manganic acid is at once decomposed 
into permanganic acid and manganese dioxide: 

3H 2 Mn0 4 = 2HMn0 4 + Mn0 2 + 2H 2 0. 

The change of a manganate to a permanganate is effected 
simply by passing carbon dioxide into the solution, or by 
boiling or allowing the solution to stand in the air. The 
change by means of carbon dioxide is represented by the 
equation 

3K 2 MnO -f 2C0 2 = 2K 2 C0 3 + Mn0 2 + 2KMn0 4 . 

With water the change takes place thus : 

3K 2 Mn0 4 + 2H ? = 2KMn0 4 + Mn0 2 + 4KOH. 

The potassium hydroxide and the manganese dioxide 
react upon each other to form a manganite of more or less 
complicated composition. While the manganates are 
decomposed by acids, forming permanganates, the latter 
are decomposed by alkalies, forming manganates. Thus 
when a solution of potassium permanganate is boiled with 
potassium hydroxide the color changes to green, owing to 
the formation of the manganate : 

2KOH + 2KMn0 4 = 2K 2 Mn0 4 + H 2 + 0. 

This change takes place readily in the presence of sub- 
stances that have the power to take up oxygen; but if such 
substances are present the reduction goes further, forming 
finally a manganite which is a derivative of the hydroxide, 
MnO(OH) 2 . 

Permanganic Acid and the Permanganates. — The sim- 
plest method of obtaining the permanganates is by decom- 
position of the manganates, as described in the last 
paragraph. 



POTASSIUM PERMANGANATE. 617 

Potassium Permanganate, KMn0 4 , is manufactured 
on the large scale by oxidizing manganese dioxide in the 
presence of a base. Sometimes the oxidation is effected 
by the oxygen of the air; sometimes by the action of an 
oxidizing agent, as potassium chlorate or nitrate. The 
fundamental reaction in each case is that represented by 
the equation 

MnO a + 2KOII + = K 2 Mn0 4 + H a O. 

As will be observed, it is a reaction of the same kind as 
that involved in the conversion of a sulphite into a sul- 
phate. Probably the first action of the hydroxide upon 
the dioxide consists in the formation of the manganite, 
K 2 MnO s , and this is then oxidized to the manganate. 
When the solution of the manganate is treated with sul- 
phuric acid a change similar to those referred to above 
takes place, and the permanganate is formed. The salt 
is easily soluble in water, and is deposited from its solution 
in crystals, isomorphous with potassium perchlorate, which 
appear nearly black, with a greenish lustre. Its solution 
in water has a purple or reddish-purple color, according 
to the concentration. Very concentrated solutions appear 
almost black. The salt is used extensively in the laboratory 
and in the arts as an oxidizing agent. Its action will be 
readily understood from what has already been said in 
regard to the conduct of manganese towards acids and 
towards alkalies. When the permanganate undergoes 
decomposition in the presence of an acid the manganese 
tends to form a manganous salt, and all the oxygen present 
in excess of that needed for this purpose is given off. 
Thus the decomposition with sulphuric acid takes place as 
represented in the equation 

2KMn0 4 + 3H 2 S0 4 = 2MnSO, + K 2 S0 4 + 3H 2 + 50. 

Therefore, when potassium permanganate is used as an 
oxidizing agent in acid solution, two molecules of the salt 
KMn0 4 give five atoms of oxygen. On the other hand, 
when the action takes place in alkaline solution the action 



618 COLLEGE CHEMISTRY. 

reaches its limit in a manganite, which, for purposes of 
calculation, may be regarded as having the composition 
K 2 Mn0 3 . The first change is from the permanganate to 
the manganate as represented in the equation 

2KMn0 4 + 2KOH = 2K 2 Mn0 4 + H 2 + 0, 

and then the manganate loses another atom of oxygen, 

2K 2 Mn0 4 = 2K 2 Mn0 3 + 20. 

Therefore, when the permanganate is used as an oxidizing 
agent in alkaline solution, two molecules of the salt yield 
three atoms of oxygen. 

The permanganates and manganates are valuable dis- 
infecting agents, and the sodium salts are extensively used 
for this purpose, under the name of Gondy's liquid. 

When a solution of barium permanganate is treated with 
sulphuric acid, free permanganic acid is obtained in solu- 
tion. It is more readily obtained by the action of an 
electric current on a solution of potassium permanganate. 
If the negative electrode is placed in a porous cup filled 
with water, and this cup in a solution of potassium per- 
manganate into which is inserted the positive electrode, 
the alkali accumulates in the porous cup and can be 
removed, and a solution of the acid remains in the outer 
vessel. The acid is unstable. When dry potassium per- 
manganate is added to concentrated sulphuric acid, oily 
drops separate and collect upon the bottom of the vessel. 
These are manganese heptoxide, Mn 2 7 , which is formed 
thus: 

2KMn0 4 + H 2 S0 4 = K 2 S0 4 + Mn 2 Y + H 2 0. 

The compound. bears to permanganic acid the relation of 
an anhydride : 

2HMn0 4 = Mn 2 7 + H 2 0. 

It is very explosive. 
Reactions which are of Special Value in Chemical 

Analysis. — The conduct of manganous salts towards solu- 
ble hydroxides has been described. The hydroxide is 



EXPERIMENTS WITH MANGANESE. 619 

Boluble in ammonia and ammoninm salts, but tliis solu- 
tion turns brown when exposed to the air and the man- 
ganese is gradually precipitated as the hydroxide Mn(OII) 3 . 

Soluble carbonates precipitate manganous carbonate. 

The conduct towards ammonium sulphide has been 
described. When oxidizing agents like hypocldoriles, 
chlorine, or bromine act upon manganous salts in solutions 
in presence of soluble hydroxides, hydroxides correspond- 
ing to the dioxide Mn() 2 , such as Mn(011) 4 , MnO(01I) 2 , 
are precipitated. Instead of the above-mentioned oxidiz- 
ing agents, potassium permanganate may be used. 

The action of potassium permanganate and manganate 
as oxidizing agents when used in alkaline and in acid 
solutions has been described above. 

Manganese is easily detected by heating the substance 
under examination w T ith nitric acid and lead peroxide, 
when permanganic acid will be formed if manganese is 
present, and its formation will be shown by the purple 
color of the solution. 

With microcosmic salt and borax manganese gives an 
amethyst-colored bead in the oxidizing flame, which 
becomes colorless in the reducing flame. 

EXPERIMENTS. 

Manganese and its Compounds. 

Experiment 237. — Make and crystallize some manganous chlo- 
ride by treating manganese dioxide with hydrochloric acid. Also 
make some manganous sulphate by heating manganese dioxide 
with sulphuric acid. Use these solutions for the purpose of 
studying the conduct of manganous salts. 

Experiment 238. — In a small porcelain crucible heat together 
5 grams manganese dioxide, 5 grams solid potassium hydroxide, 
and 2| grams potassium chlorate. When the mass has turned 
green, dissolve the contents in water and neutralize' most of the 
free alkali in the solution. Or pass carbon dioxide through the 
solution without boiling. 



CHAPTER XXXIII. 

ELEMENTS OF FAMILY VIII, SUB-GROUP A. 
IRON.— COBALT.— NICKEL. 

General. — The three elements which form this group 
are in many respects very similar, and their atomic weights 
differ but little from one another. That of iron (56) is 
nearly the same as that of manganese (55), while cobalt 
and nickel have nearly the same atomic weight. There is 
much in iron that suggests manganese. It forms two 
series of compounds, the ferrous and ferric compounds, 
which are analogous to the manganous and manganic com- 
pounds. In the first series iron appears to be bivalent, as 
shown in the formulas 

FeCl 2 , Fe(OH) 2 , FeO, FeS, FeS0 4 , FeC0 3 , etc. 

In the second series it appears to be trivalent, as indicated 
in the formulas 

FeCl 3 , Fe(OH) 3 , Fe 2 3 , Fe(N0 3 ) 3 , Fe 2 (SOJ 3 , etc. 

Like chromium and manganese it also forms an acid 
known as ferric acid, H 2 Fe0 4 , which in composition is 
analogous to chromic and manganic acids. The soluble 
salts of this acid are, however, unstable, and on decompos- 
ing yield ferric hydroxide. Oxidizing agents readily con- 
vert ferrous compounds into ferric compounds, and 
reducing agents reconvert the latter into the former. 
When exposed to the air most ferrous compounds are 
oxidized to ferric compounds. The ferrous compounds 
in which iron is bivalent are similar to the compounds of 
the zinc group. The ferric compounds, however, in which 
the iron is trivalent, are similar to the aluminium com- 
pounds; and in ferric acid it exhibits a resemblance to 

620 



IRON. 621 

chromium. Cobalt and nickel resemble iron in respect to 
their power to form two series of compounds correspond- 
ing to the ferrous and ferric compounds. Both elements 
preferably form compounds of the lower series, examples 
of which are represented by the formulas 

CoCl 2 , Co(OH) 2 , CoO, Co(N0 3 ) 2 , CoS0 4 , etc. 
NiCl 8 , Ni(OH) 2 , NiO, Ni(N0 3 ) 2 , NiS0 4 , etc. 

Cobalt forms a few compounds corresponding to the 
ferric series ; and nickel forms a hydroxide, of the formula 
Ni(OH) 3 . While the power of cobalt to form compounds 
in which it is trivalent is much weaker than that of iron, 
it is stronger than that of nickel, the latter being almost 
exclusively bivalent. 

Iron, Fe (At. Wt. 56). 

Introductory.— The importance of this metal to man- 
kind can hardly be overestimated, and for many centuries 
it has played a commanding part in the industries. It 
requires little thought to convince one that without it the 
earth would be quite a different place from what it now is. 
In the earliest periods of history metals were but little 
used, as but few of them are furnished ready for use by 
nature. Stones were therefore first used, and these were 
shaped into a variety of implements, many of which still 
exist, and furnish evidence of the Stone Age. After a 
time copper and tin were used in the form of an alloy or 
bronze, as copper is found in nature in the free condition. 
During this period, known as the Bronze Age, stone im- 
plements gave way to those made of bronze. Afterwards 
men learned to extract iron from its ores, and the Iron 
Age was introduced; and this has continued up to the 
present, as nothing has since been found which can 
advantageously take the place of iron. 

Forms in which Iron occurs in Nature, — Iron occurs in 
small quantity native in meteorites, in the basalts of 
Bohemia and Greenland, and in some gabbros. The iron 
meteorites always contain nickel, and frequently small 



622 



COLLEGE CHEMISTRY. 



quantities of other elements, as manganese and carbon. 
Compounds of iron occur in enormous quantities, and 
widely distributed in the earth. Among the more impor- 
tant are the following-named : hematite, Fe 2 3 ; magnetite, 
Fe 3 4 ; brown iron ore, Fe 4 3 (OH) 6 ; siderite, or the car- 
bonate, FeC0 3 ; pyrite, FeS 2 ; pyrrhotite, Fe 7 S 8 . It is also 
contained in many silicates in small quantity, and in con- 
sequence of the disintegration of the constituents of rocks 
it is found in the soil, and in many natural waters. In 
the vegetable kingdom it is always found in chlorophyll, 
and in the animal kingdom always in the blood. The 
compounds which are chiefly used for the purpose of 
making iron, or the iron ores, are magnetite, Fe s 4 ; 
hematite, Fe 2 3 ; brown iron ore, Fe 4 3 (OH) 6 ; and spathic 
iron, or siderite, FeC0 3 . 

Metallurgy,— The ores of iron, after they are broken 



up, are 



first roasted, in order to drive off water from the 
hydroxides; to decompose car- 
bonates; to oxidize sulphides; and, 
as far as possible, to convert the 
oxides into ferric oxide, Fe 2 3 , 
which is the most easily reducible 
of the oxides of iron. After the 
ores are prepared in this way they 
are reduced by heating them with 
carbon and fluxes in the blast- 
furnaces, when the iron collects 
in the molten condition under the 
so-called slag at the bottom of the 
f urnace. Blast - furnaces differ 
somewhat in construction, but the 
essential parts are represented in 
Fig. 78. 

The inner cavity of the furnace 
is narrow at the top and bottom, 
as is shown in the figure. Through 
pipes, known as tuyeres, such as 
that represented at the lower part of the left-hand side of 




Fig. 78. 



IRON. 623 

the figure, air is blown into the furnace to facilitate the 
combustion. In modern furnaces arrangements are made 
:'1)<)\ e for carrying oil' the gases and utilizing them as fuel. 
The inner walls are built of fire-bricks, and these are sur- 
rounded by ordinary bricks, or stone-work. The furnaces 
vary in height from 25 to 80 or 90 feet, an average height 
being about 45 feet. The reduction of the ores is accom- 
plished by placing in the furnace alternating layers of coke 
or charcoal, and the ores mixed with proper fluxes. The 
nature of the flux depends upon the ore. If this contains 
silicon dioxide or clay, lime is added; while, if it contains 
considerable lime, minerals rich in silicic acid are used, 
such as feldspar, clay-slate, etc. The object of the flux is 
to form a slag in which the reduced iron collects, and by 
which it is protected from oxidation. When the fire is 
once started in a blast-furnace the operation of reduction 
is continuous until the furnace is burned out. Alternate 
layers of ore and flux and carbon are added, and, as the 
reduced iron collects below, it is from time to time draw T n 
off and allowed to solidify in moulds of sand. The reduc- 
tion is largely accomplished by carbon monoxide.- In the 
lower part of the furnace the fuel burns to carbon dioxide, 
but this comes in contact with hot carbon, and is then 
reduced to the monoxide. The hot monoxide in contact 
with the oxides of iron reduces these, and is itself con- 
verted into the dioxide. A large proportion of the carbon 
monoxide, however, escapes oxidation, and this is carried 
off from the top of the furnace to the bottom by properly 
arranged pipes, and is then utilized as fuel. 

Varieties of Iron. — The iron obtained as above described 
is known as pig iron or cast iron. It is quite impure, 
containing carbon, phosphorus, sulphur, silicon, etc. If, 
when drawn from the furnace, the iron is cooled rapidly, 
nearly all the carbon contained in it remains in chemical 
combination, and the iron has a silver-white color. This 
product is known as white cast iron. If the iron cools 
slowly, most of the carbon separates as graphite, and this 
being distributed through the mass gives it a gray color. 



624 COLLEGE CHEMISTRY. 

This product is known as gray cast iron. If the ore con- 
tains considerable manganese, this is reduced with the 
iron, and iron made from such ores and containing man- 
ganese has the power to take up more carbon than ordinary 
iron. This product, containing from 3.5 to 6 per cent 
combined carbon, is known as spiegel-iron. 

All varieties of cast iron are brittle and easily fusible. 
The gray iron fuses at a lower temperature than the white, 
and is not as brittle; it is therefore well adapted to making 
castings. When cast iron is treated with hydrochloric 
acid the carbon which is present in combined form is given 
off in combination with hydrogen as hydrocarbons, some 
of which, have a disagreeable odor. This is, of course, the 
cause of the bad odor noticed in dissolving ordinary cast 
iron in acids. The uncombined or graphitic carbon, on 
the other hand, remains undissolved. Owing to its brittle- 
ness, cast iron cannot be welded. When the carbon, sili- 
con, and phosphorus are removed the iron becomes tough 
and malleable, and its melting-point is much raised. The 
product thus obtained is known as wrought iron. 

Puddling. — Wrought iron is obtained, from cast iron by 
the puddling process. The puddling furnace has a flat 
oval hearth and low arched roof. The sides of the hearth 
are lined with a layer of iron ore (oxide). Coal is burned 
on a grate and the flame passes into the furnace at one end 
and out at the other, thus coming in contact with the roof 
and the charge of iron. By contact with the flame, and 
by the heat radiated from the roof, the cast iron melts. 
The carbon and silicon are removed from the molten cast 
iron, partly by the oxygen in the air or flame, but princi- 
pally by the oxygen in the iron ore, which is itself thus 
reduced to wrought iron. 

Wrought iron contains less than 0.6 per cent of carbon, 
and, as the percentage of carbon decreases, the malleability 
increases and the melting-point rises. The melting-point 
of good wrought iron is from 1900° to 2100°. Small 
quantities of sulphur, phosphorus, silicon, and manganese 
exert a very marked influence upon its properties. The 



IRON. 625 

process of welding consists in heating two pieces of iron to 
a high temperature, putting some borax upon one of them, 
laving them together, and hammering, when, as is well 
known, they adhere firmly together. The object of the 
borax is to keep the surfaces bright, which it does by 
uniting with the oxide and forming an easily fusible 
borate. 

Bessemer Process. — Molten cast-iron is poured into a 
large vessel called the converter. The carbon and silicon 
are entirely oxidized and removed by means of a blast of 
air forced through the metal from below. No fuel is used, 
as the heat generated by the oxidation of carbon and 
silicon is sufficient to raise the temperature above 2100°. 
The converter contains molten wrought iron after the 
oxidation. By addition of spiegel-iron a product contain- 
ing any desired percentage of carbon is obtained. 

Thomas- GilcTirist Process. — Iron which contains more 
than a very small percentage of phosphorus is not adapted 
to the manufacture of Bessemer steel in the ordinary way; 
but it has been found that, if the converters are lined with 
lime and magnesia, such iron may be used. Under these 
circumstances the phosphorus is oxidized, and forms 
calcium and magnesium phosphates, which are of value as 
fertilizers (see Calcium Phosphate). This process is known 
as the Thomas- Gilchrist or the basic-lining process. 

Siemens- Martin Furnace. — This is simply a furnace in 
Avhich gas is used as fuel, the gas being previously heated 
in a Siemens regenerative furnace. This is used in making 
steel by heating together pig iron and pure wrought iron. 

Steel and Wronght Iron. — The product of the puddling 
furnace is called wrought iron; while those formed in the 
Bessemer process and in the Siemens- Martin furnace are 
called steel. Bessemer steel often contains less than 0.6 
per cent of carbon, and Siemens-Martin steel is the purest 
form of wrought iron, containing less carbon and silicon 
than the product of the puddling furnace. 

Tempering. — When steel is heated and cooled suddenly, 
it is rendered extremely hard and brittle; and when 



626 COLLEGE CHEMISTRY. 

hardened steel is carefully heated, and allowed to cool 
slowly, it becomes very elastic. This process is called 
tempering. 

Properties of Iron. — Pure iron is almost unknown. Of 
the commercial varieties, it follows from what has been 
said that wrought iron is the purest. That which is used 
for piano-strings is the purest iron obtainable in the 
market; it contains only about 0.3 per cent of impurities. 
Pure iron can be made in the laboratory by igniting the 
oxide or oxalate in a current of hydrogen, and by reducing 
ferrous chloride in hydrogen. In larger quantity it can 
be prepared by melting the purest wrought iron in a lime 
crucible by means of the oxyhydrogen flame. The im- 
purities are taken up by the crucible, and a regulus of the 
pure metal is left behind. That made by reduction of the 
oxide or oxalate is, of course, in finely divided condition. 
If in its preparation the temperature is kept as low as 
}30ssible, the product takes fire when brought in contact 
with the air: while if the temperature is high, the product 
has not this power. Iron is white, and is one of the 
hardest metals: and its melting-point is higher than that 
of wrought iron. Pure iron is attracted by the magnet. 
In contact with a magnet, or when placed in a coil through 
which an electric current is passing, it becomes a magnet; 
but the purer it is the sooner it loses the magnetic power 
when removed from the magnet or the coil. Steel, how- 
ever, retains its magnetism. "When heated to a sufficiently 
high temperature iron burns, and forms the oxide, Fe 3 4 . 
This takes place much more easily in oxygen than in the 
air. In dry air iron does not undergo change, but in moist 
air it rusts, or it becomes covered with a layer of oxide 
and hydroxide, which is formed by the action of the air, 
carbon dioxide, and water. Water that contains salts in 
solution facilitates the rusting. Various methods are 
adopted to protect iron from this change, most of which 
are, however, purely mechanical. A method which 
promises valuable results is that invented by Barff, which 
consists in introducing the iron into water vapor at a 



FERROUS CHLORIDE FERRIC CHLORIDE. 627 

temperature of 650°, when it becomes covered with a, 
firmly adhering layer of oxide. 

Iron dissolves in acids with evolution of hydrogen, and 
generally with formation of ferrous salts: 

Fe + 2HC1 = FeCi, + II a ; 
Fe + H a S0 4 = FeS0 4 + H 2 . 

When cold nitric acid is used, ferrous nitrate and am- 
monium nitrate are the products; if the acid is warmed, 
ferric nitrate and oxides of nitrogen are formed. When 
an iron wire which has been carefully polished is intro- 
duced for an instant into red fuming nitric acid it can 
afterward be put into ordinary nitric acid without under- 
going change. It is said to be in the passive state; and 
the commonly accepted explanation of the phenomenon 
is that the wire is covered with a thin layer of oxide. 
As, however, the passive condition is lost by contact with 
an ordinary wire, the explanation does not appear to be 
adequate. 

Ferrous Chloride, FeCl 2 .— This is a colorless mass, 
which deliquesces in the air, is volatile at a high tempera- 
ture, and determinations of the specific gravity of its 
vapor made at very high temperatures have shown that its 
molecule under these conditions should be represented by 
the formula FeCl 2 . At lower temperatures the molecule 
appears to be more complex. The evidence on this point 
is not conclusive. 

A solution of ferrous chloride made by dissolving iron 
in hydrochloric acid is used in medicine under the name 
Liquor Ferri chlorati. It contains 10 per cent iron. 

Ferric Chloride, FeCl 3 .— Ferrous chloride is readily 
converted into ferric chloride by oxidation. The simplest 
way to make a solution of the ferric compound is to dis- 
solve iron in hydrochloric acid and pass chlorine into it 
to complete saturation. Anhydrous ferric chloride is 
obtained by heating iron wire in dry chlorine. It forms 
black, lustrous crystalline laminae, is volatile at a lower 



628 . COLLEGE CHEMISTRY. 

temperature than the ferrous compound, and the specific 
gravity of the vapor is that required by a compound whose 
molecule is represented by the formula FeCl 3 . When 
treated with nascent hydrogen, ferric chloride is converted 
into ferrous chloride: 

FeCl 3 + H = FeCl 2 + HC1. 

A solution of ferric chloride is used in medicine under the 
name Liquor Ferri sesquichlorati. 

Cyanides. — -The compounds which iron forms with 
cyanogen are of special interest. The simple compounds, 
ferrous cyanide, Fe(CN) 2 , and ferric cyanide, Fe(CN) 3 , 
corresponding to the above-mentioned chlorides, are not 
known, but double compounds of these with other cyanides 
are well known, and some of them are manufactured on 
the large scale. When a solution of potassium cyanide 
acts upon metallic iron or the oxides of iron, a solution is 
formed from which the salt known as potassium ferro- 
cyanide or yellow prussiate of potash crystallizes. This 
has the composition K 4 Fe(CN) 6 + 3H 2 0, and may be 
regarded as made up of a molecule of ferrous cyanide and 
four molecules of potassium cyanide, as represented in 
the formula Fe(CN) 2 .4KCN + 3H 2 0. When this salt is 
treated with chlorine it is converted into potassium ferri- 
cyanide, or red prussiate of potash, K 3 Fe(CN) 6 , which is 
to be regarded as consisting of ferric cyanide and potassium 
cyanide, as represented in the formula Fe(CN) 3 .3KCN. 
The transformation is represented thus : 

K 4 Fe(CN) 6 + CI = K 3 Fe(CN) 6 + KC1. 

From these two a number of other cyanogen compounds 
are obtained. When treated in concentrated solution with 
concentrated Irydrochloric acid they yield the free acids, 
and by treating them with solutions of different metallic 
salts corresponding salts of these acids are obtained. 
Among the most important of these derivatives are the 
following : 



POTASSIUM FERROCYANIDE, ETC. 621) 

Ferrohydrocyanic acid, Ferriliydrocyanic acid, 

H i Fe(CN) a H 8 Fe(CN) fl 

Potassium ferrocyanide, Potassium ferricyanide, 

K 4 Fe(CN) a K,Fe(CN) 8 

Sodium ferrocyanide, Sodium fei ricyanide, 

Na 4 Fe(CN) a Na a Fe(CN)a 

Barium ferrocyanide, 

Ba a Fe(CN)« 

Ferric ferrocyanide Ferrous ferricyanide, 

Fe 4 [Fe(C>s T ) 6 ]3 Fe, [Fe(CN),] a 

Ferri-potassium ferrocyanide, 

KFeFe(CN) a 

Potassium Ferrocyanide, K 4 Fe(CN) 6 + 3H 2 0. — As 
stated above, this salt can be made by treating iron or 
the oxides of iron with a solution of potassium cyanide. 
On the large scale it is manufactured by melting crude 
potash or potassium carbonate, and gradually adding 
a mixture of iron filings or turnings, and refuse animal 
matter, as claws, horns, hoofs, hair, etc. Or the potash 
is melted with the animal substances and potassium 
cyanide thus formed, and this treated in solution with 
ferrous carbonate, when the ferrocyanide is formed. It 
forms large yellow pyramids belonging to the tetragonal 
system. At the ordinary temperature it dissolves in three 
to four parts of water, and more easily in hot water. It 
gives up its water of crystallization very easily. When 
heated it is decomposed, forming potassium cyanide, 
nitrogen, and a compound of iron and carbon : 

K 4 Fe(CN) 6 = 4KCN + N 2 + FeC 2 . 

Ferrohydrocyanic Acid, H 4 Fe(CN) 6 , formed as above 
described, is a white crystalline substance, which is easily 
soluble in water and alcohol. The ferric salt is the sub- 
stance commonly called insoluble Prussian blue. The 
relation of the salt to the acid is shown by the formulas 

H 4 Fe(CN)' 6 Fe t [Fe(CN)J, 

Ferrohydrocyanic acid Ferric ferrocyanide, or 

Prussian blue 



630 COLLEGE CHEMISTRY. 

Ferric Ferrocyanide, or Prussian Blue, Fe 4 [Fe(CN)J s . 
— This compound is readily formed by adding a solution 
of a ferric salt to a solution of potassium ferrocyanide, and 
appears as a dark-blue precipitate: 

3K 1 Fe(CN) 6 + 4FeCl 3 = Fe 4 [Fe(CN) 6 ] 3 + 12KC1. 

Potassium Ferricyanide, K 3 Fe(CN) 6 . — This salt is 
formed by treating the ferrocyanide, either dry or in solu- 
tion, with chlorine. It forms large, dark-red, monoclinic 
prisms. It dissolves in about three times its weight of 
waier at the ordinary temperature, and is more easily 
soluble in hot water. 

Ferrihydrocyanic Acid, H 3 Fe(CN) 6 , is a crystallized 
substance. 

Ferrous Ferricyanide, Fe 3 [Fe(CN) 6 ] 2 , is commonly 
called TurnbulPs blue. It is formed by adding potassium 
ferricyanide to a solution of ferrous sulphate, or any 
ferrous salt: 

3FeS0 4 + 2K 3 Fe(CN) 6 = Fe 3 [Fe(CN)J 2 + 3K 2 S0 4 . 

Ferrous Hydroxide, Fe(OH)., , is formed when a soluble 
hydroxide is added to a solution of a ferrous salt. It is a 
white precipitate, but it is usually obtained as a greenish 
mass, as it is very easily oxidized by the oxygen of the air 
and that contained in the solutions. When allowed to 
stand in contact with the air it turns dirty green, and 
finally brown, being converted into ferric hydroxide. 
When heated in the air it loses water, and takes up 
oxygen, forming ferric oxide. 

'Ferric Hydroxide, Fe(OH) 3 . — This compound is formed 
most readily by addiug ammonia to a solution of a ferric 
salt, when it appears as a voluminous brownish-red pre- 
cipitate. When filtered, washed, and dried, its composi- 
tion is not changed. If heated to .100°, or if the solution 
is boiled for some time, it loses water, and forms com- 
pounds of the formulas FeO.OH, Fe 2 0(OH)„ etc. The 



FERRIC HYDROXIDE, ETC. 63 1 

Latter is derived from the normal hydroxide as represented 
in the equation 

2Fe(OH) 3 = Fe 2 0(OII) 1 + H 2 0. 

The mineral pyrosiderite is the hydroxide FeO.OH. 

Brown iron ore is Fe 4 8 (OH) 6 ; and bog iron ore is 
Fe 2 0(OII) 4 . All of these are derivatives of the normal 
hydroxide. The normal hydroxide differs from aluminium 
hydroxide in the fact that it has no acid properties. 
Therefore, if the two hydroxides are treated together with 
a caustic alkali only the aluminium hydroxide dissolves. 
The compound FeO.OH, corresponding to AIO.OH and 
CrO.OII, yields salts under some circumstances. Thus a 

calcium salt, t^c\ r\>^> is formed by heating together 

ferric oxide and lime to a high temperature. In composi- 
tion this is plainly analogous to the spinels. Magnetic 
oxide of iron or magnetite is believed to be the correspond- 
ing ferrous salt, -^ Q*/-v>Fe. Franklinite also is a salt of 

the same order, containing zinc. It is essentially a zinc 

FeO 
salt, of the formula y q'q>Zii, but some of the zinc is 

replaced by iron and manganese. 

Ferrous-Ferric Oxide, Fe 3 4 . — As stated above, this 
compound is regarded as analogous to the spinels, and as 
the ferrous salt of the acidic hydroxide FeO.OH, as repre- 
sented in the equation (FeO.O) 2 Fe. It is found in nature 
as the mineral magnetite, and loadstone, which occurs in 
'Sweden, Norway, and elsewhere. It is, further, formed 
when iron is burned in oxygen, and when water is passed 
over red-hot iron. Some of the magnetite which occurs 
in nature has the power to attract iron, or is magnetic. 

Soluble Ferric Hydroxide is formed when a -solution of 
ferric chloride or ferric acetate is treated with ferric 
hydroxide, and the solution thus formed dialyzed (see 
p. 395). The ferric salts pass through the membrane, 
and the ferric hydroxide remains in solution in water, 



632 COLLEGE CHEMISTRY. 

forming a deep-red liquid. It is used in medicine. Small 
quantities of salts cause the precipitation of ferric hy- 
droxide from the solution. 

Ferric Oxide, Fe 2 3 , is found in nature, and is known 
as hematite, forming one of the most valuable ores of iron. 
It can be made in the laboratory by igniting the hydroxide. 
As hematite, it is a black, crystallized substance with a 
high lustre. Otherwise it has a red or a reddish-brown 
color. The oxide found in nature and that which has 
been strongly ignited are very difficultly soluble in acids. 
In the preparation of fuming sulphuric acid by heating 
ferrous sulphate (see p. 232) there is left a residue of ferric 
oxide known as rouge, which is used as a red pigment and 
as a polishing powder. A specially fine variety of rouge 
for polishing is manufactured by heating ferrous oxalate, 
FeC 2 4 , in contact with the air. 

Ferrous Sulphide, FeS, is formed by direct union of 
iron and sulphur when the two are heated together. It 
is manufactured by heating iron filings and sulphur 
together in a crucible. The pure compound is yellow and 
crystalline. When heated in contact with the air it is 
oxidized to ferrous sulphate, if the temperature is not too 
high. At a higher temperature the products are sulphur 
dioxide and ferric oxide. When a solution of a ferrous 
salt is treated with ammonium sulphide, ferrous sulphide 
is precipitated as a black powder. When a ferric salt is 
treated with ammonium sulphide it is reduced to the 
ferrous condition, and then ferrous sulphide is precipi- 
tated : 

Fe 2 (S0 4 ) 3 + (NH 4 ) 2 S = 2FeS0 4 + (NH 4 ) 2 S0 4 -f S; 
2FeS0 4 + 2(NHJ 2 S = 2FeS + 2(NHJ,S0 4 . 

The sulphide thus obtained is readily oxidized in the air, 
and forms the sulphate. The compact variety is used in 
making hydrogen sulphide (which see). 

Ferrous Carbonate, FeC0 3 . — This salt occurs in nature 
as spathic iron or siderite. It crystallizes in forms similar 



FERROUS SULPHATE, ETC. 633 

bo those of calc spar or. calcium carbonate, CaC0 8 . Like 
this, further, it dissolves in water which contains carbon 
dioxide, and is therefore contained in natural waters which 
come in contact with it. When a solution of a ferrous salt 
is treated with a soluble carbonate a white precipitate is 
formed, which is ferrous carbonate; but in contact with 
the air this is rapidly oxidized and decomposed, leaving 
ferric hydroxide, which with carbonic acid does not form 
a salt. In this respect ferric hydroxide acts like aluminic 
and chromic hydroxides, and therefore when a soluble 
carbonate is added to a solution of a ferric salt the 
hydroxide and not ferric carbonate is thrown down. 

Ferrous Sulphate, FeS0 4 . — This important compound 
is manufactured on the large scale by the spontaneous 
oxidation of pyrite in contact with the air, and by dissolv- 
ing iron in sulphuric acid. It is frequently called " green 
vitriol" (see p. 544), and more commonly "copperas." 
Under ordinary conditions it crystallizes in transparent, 
green, monoclinic crystals with seven molecules of water, 
just as zinc sulphate, magnesium sulphate, etc., do; and 
when heated, six of these are given off readily, while the 
last is given off with difficulty — a fact which makes it 
appear probable that the salt is a derivative of tetra- 
hydroxyl-sulphuric acid, as represented in the formula 

OS < p . While it ordinarily crystallizes in monoclinic 

crystals, it takes the rhombic form if its supersaturated 
solution is touched with a crystal of zinc sulphate. It 
also crystallizes in the triclinic system with five molecules 
of water, like cupric sulphate, if a crystal of the latter salt 
is placed in its concentrated solution. 

Ferric Sulphate, Fe 2 (SOJ 3 .— This salt is formed by 
oxidation of ferrous sulphate. It is also formed by dis- 
solving ferric oxide or hydroxide in sulphuric acid. 

Ferrous Phosphate, Fe 3 (P0 4 ) 2 , occurs in nature crystal- 
lized with eight molecules of water as the mineral vivianite. 
Both this salt and ferric phosphate, FeP0 4 , are insoluble 



634 COLLEGE CHEMISTRY. 

and are formed when solutions of ferrous and ferric salts 
are treated with sodium phosphate. 

Iron Disulphide, FeS 2 , is not analogous to any oxygen 
compound of iron. In it the metal appears to be quad- 
rivalent. The disulphide occurs very widely distributed 
and in large quantities in nature as the mineral iron pyrites 
or pyrite, which crystallizes in the regular system, and as 
marcasite, which crystallizes in the rhombic system. It 
can be made artificially, and if crystallized it appears in 
the form of pyrite. Its conduct under the influence of 
heat has been repeatedly referred to in connection with 
the roasting of iron and other ores. As pyrite it has a 
golden-yellow color, and it has frequently been taken for 
gold by those not familiar with it. The name "fool's 
gold," by which it is sometimes popularly known, suggests 
this fact. 

Arsenopyrite, FeAsS, also called arsenical pyrites, occurs 
in nature, and is a valuable source of the element arsenic: 
for, as has been stated (see p. 305), when it is heated it 
gives off arsenic, and ferrous sulphide is left behind. 

Reactions which are of Special Value in Chemical 
Analysis. — Ferrous Compounds. — The reactions of ferrous 
compounds with the soluble hydroxides and carbonates, 
ammonium sulphide, potassium ferricyanide, and with 
oxidizing agents have been explained above. With am- 
monium salts ferrous chloride forms double salts, which 
are soluble; therefore, if ammonium chloride is added to 
a solution of the salt ammonia does not precipitate the 
hydroxide. Further, ammonia does not completely pre- 
cipitate the hydroxide from a solution of a ferrous salt, as 
an ammonium salt is formed. By standing in the air, 
however, these solutions containing the double salts are 
oxidized, and ferric hydroxide is precipitated. The reac- 
tions with potassium cyanide will be understood from what 
has been said concerning the compounds of ferrohydro- 
cyanic and ferrihydrocyanic acids. 

Ferric Compounds. — The reactions of ferric compounds 
with the soluble hydroxides and carbonates, ammonium 



COBALT. 635 

sulphide, potassium ferrocyanide, and potassium ferri- 
cyanide have been explained above. When hydrogen 
sulphide is passed through a solution of a ferric salt, 
reduction to the corresponding ferrous suit takes place, 
and sulphur separates, which gives the solution a milky 
appearance: 

Fe 2 (S0 4 ) 3 + II 2 S --= 2FeS0 4 + II 2 S0 4 + S; 
2FeCl s + H a S = 2FeCl 3 + 2HC1 -f S. 

When a neutral solution of a ferric salt is treated with 
suspended barium carbonate the iron is precipitated as the 
hydroxide. 

When a neutral solution of a ferric salt is treated with 
acetate of potassium or sodium it turns dark red, in conse- 
quence of the formation of ferric acetate which remains 
in solution. When the solution is boiled the acetate 
breaks down into acetic acid and a basic hydroxide of iron, 
which is precipitated. When the basic acetate is filtered 
off, washed, and heated it is changed to ferric hydroxide: 

FeCl 3 + 3NaC 2 H 3 2 = Fe(C a H 8 O a ), + 3NaCl; 
Fe(C 2 H 3 O a ) 3 + 3II 2 = Fe(OH) 3 + 3C 2 H 4 2 . 

When potassium sulphocyanate, KCNS, is added to a solu- 
tion of a ferric salt a blood-red color is produced. This 
occurs even in extremely dilute solutions of ferric salts. 

The borax bead is colored bottle-green in the reducing 
flame, and brown-red to yellowish red in the oxidizing 
flame, when treated with compounds of iron. 



Cobalt, Co (At. Wt. 59). 

General. — As stated in the remarks introductory to the 
iron group, cobalt, like nickel, preferably forms com- 
pounds which are analogous to ferrous compounds. It, 
however, forms a few which are analogous to ferric com- 
pounds, its power in this direction being greater than that 



636 COLLEGE CHEMISTRY. 

of nickel. Its salts form a great variety of compounds 
with ammonia, and these have been extensively studied. 

Occurrence and Preparation. — Cobalt occurs in nature, 
almost always in company with nickel. The principal 
minerals containing it are smaltite, CoAs 2 , and cobaltite, 
CoS 2 .CoAs 2 . In each of these iron, and generally some 
nickel, take the place of a part of the cobalt. 

Properties. — Cobalt has a silver-white color, with a 
slight cast of red. It is harder than iron, and melts at a 
somewhat lower temperature; is tenacious; and has the 
specific gravity 8.9. It dissolves in nitric acid. 

Cobaltous Chloride, CoCl 2 , is formed by heating cobalt 
in chlorine gas, and in solution by treating cobalt car- 
bonate with hydrochloric acid. From the solution it 
crystallizes in dark-red prisms of the composition CoCl 2 
-f- 6H 2 0. The anhydrous salt is blue. When the blue 
salt is treated with water it turns red, and when the red 
salt is heated it turns blue. This difference in color 
between the anhydrous and the hydrated salts is charac- 
teristic of cobalt salts. If marks are made on paper with 
a dilute solution of one of the salts the color is not per- 
ceptible. If, however, the paper is held before a fire, the 
salt loses water and turns blue, and as the blue is more 
intense than the red, it is visible. When the salt becomes 
moist again it becomes invisible. This is the basis for the 
preparation of the so-called sympathetic inks. 

Cobaltous Hydroxide, Co(OH) 2 , is formed as a red pre- 
cipitate when a soluble hydroxide is added to a cobaltous 
salt, and the blue precipitate, which is first formed and 
which is a basic salt, is allowed to stand. It is oxidized 
by contact with the air, forming cobaltic hydroxide, which 
breaks down into cobaltic oxide, Co 2 3 . 

Cobaltic Hydroxide, Co(OH) 3 , is formed when calcium 
hypochlorite is added to a solution of a cobaltous salt, and 
is a black powder. When heated it is converted into black. 
colaltic oxide, Co 2 3 . 

Cobalt Sulphide, CoS, is the black precipitate which is 
formed by adding ammonium sulphide to a cobaltous salt. 



NICKEL. 637 

It is not soluble in dilute acids, and differs from ferrous 
sulphide in this respect. Other sulphides of cobalt are 
those of the formulas Co 3 S 4 and CoS 2 . The former is 
found in nature, and is known as linnaeite. The latter 
occurs m combination with other sulphides, and with 
arsenides; as in cobaltite, CoS 2 .CoAs 2 . 

Cyanides. — Cobaltous cyanide is an insoluble dirty-red 
compound which is formed when potassium cyanide is 
added to a solution of a cobalt salt. It dissolves in an 
excess of potassium cyanide, forming a double cyanide, 
K 4 Co(CN) 6 , which is analogous to potassium ferrocyanide. 
When this solution is boiled the cyanide is oxidized, form- 
ing a compound analogous to potassium ferricyanide, thus : 

2K 4 Co(CN) 6 -f H 2 + = 2K 3 Co(CN) 6 + 2KOH. 

This acts like the corresponding iron compound. The 
cobalt is not precipitated from it by ammonium sulphide 
or sodium hydroxide. This conduct towards potassium 
cyanide distinguishes cobalt from nickel salts. 

Smalt. — The beautiful pigment known by this name is 
essentially a cobalt glass in which cobalt' takes the place 
of calcium. It is made by heating compounds of cobalt 
with quartz and potassium carbonate. The glass thus 
formed is powdered very finely and used as a pigment. It 
does not change color in the sunlight, and is not affected 
by acids nor by alkalies. 



Nickel, Ni (At. Wt. 58.7). 

General, — Nickel differs from cobalt in respect to the 
difficulty with which it forms nickelic compounds or those 
in which it is trivalent. At the same time it does form 
an oxide of the composition M 2 3 , and the corresponding 
hydroxide, Ni(OH) 3 . In all other compounds it is bivalent, 
the compounds being analogous to ferrous compounds. 

Occurrence and Preparation. — Nickel occurs native in 
meteorites. The principal minerals containing it are the 



638 college chemistry. 

arsenide, MAs, known as niccolite, and the sulpharsenide, 
MSAs or NiS 2 . MAs 2 , known as gersdormte. 

Properties. — Nickel is a white metal with a slight cast 
of yellow. It is very hard, and capable of a high polish. 
The metal in its ordinary condition is brittle, but after 
treatment with a little magnesium it becomes very mal- 
leable. Its specific gravity is 8.9, and it melts at a high 
temperature. It is not changed in the air; it dissolves 
slowly in hydrochloric and sulphuric acids, and readily in 
nitric acid. Like iron, it is magnetic. 

Alloys. — Alloys of nickel are extensively used. Argen- 
tan or German silver consists of copper, zinc, and nickel. 
Various nickel alloys are used for making coins. The 
5-cent pieces in the United States are made of an alloy 
consisting of 25 per cent nickel and 75 per cent copper. 
In Switzerland, and Belgium also, nickel coins are used. 

Other Applications of Nickel. — Besides as a constituent 
of important alloys, nickel is extensively used at present 
in nickel-plating. Iron objects are covered with a thin 
layer of the metal for the purpose of protecting them from 
rusting. The plating is accomplished as silver-plating 
and copper-plating are — by means of electrolysis, a bath 
of nickel-ammonium sulphate being used. 

Nickelous Hydroxide, M(OH) 2 , is formed when a nickel 
salt is treated with a soluble hydroxide, and is a green 
insoluble substance. When heated it is converted into 
the green oxide, NiO. 

Nickelic Hydroxide, Ni(OH) 3 , is precipitated as a black 
powder when a solution of a nickel salt is treated with 
sodium hypochlorite. 

Cyanides. — When potassium cyanide is added to a solu- 
tion of a nickel salt, nickel cyanide, M(CN) 2 , is precipi- 
tated as a greenish-white substance. With an excess of 
potassium cyanide this forms the salt, Ni(CN) 2 .2KCN, 
which, owing to the fact that nickelous salts are not con- 
verted into nickelic salts by oxidation, does not undergo 
change when boiled with potassium cyanide. When 
hydrochloric acid is added to a solution of the double 



REACTIONS OF COBALT AND NICKEL 639 

cyanide, nickelous cyanide is precipitated. If boiled with' 
precipitated mercuric oxide the double cyanide is decom- 
posed and nickel oxide is thrown down. 

Nickel Carbonyl, Ni(CO) 4 . — This interesting compound 
is formed when finely- divided nickel, such as is obtained 
by reducing nickel oxide by hydrogen at about 400°, is 
allowed to cool in a slow current of carbon monoxide. A 
gas is formed which can easily be condensed, its boiling- 
point being 43°. At — 25° it solidifies, forming needle- 
shaped crystals. When the gas is passed through a heated 
tube pure nickel is deposited. Advantage of this fact is 
taken for the purpose of preparing pure nickel on the large 
scale. Cobalt does not form a compound of this kind. 

Reactions of Cobalt and Nickel which are of Special 
Value in Chemical Analysis. — The reactions with the 
soluble hydroxides have been explained. With ammonium 
sulphide both give black sulphides, which are not easily 
dissolved by dilute hydrochloric acid. From solutions of 
the acetates hydrogen sulphide precipitates the sulphides. 
Nickel sulphide is slightly soluble in ammonium sulphide, 
and the solution has a brownish-yellow color. 

The action of the hypochlorites upon solutions of nickel 
and cobalt salts has been explained above. The reactions 
with potassium cyanide have also been explained. These 
furnish a good method for separating the two metals. 

When a solution of potassium nitrite is added to a solu- 
tion of a cobalt salt containing free acetic acid or nitric 
acid, a precipitate of cobaltic potassium nitrite is formed. 
This is a compound of cobaltic nitrite, Co(N0 2 ) 3 , and 
potassium nitrite, of the composition Co(N0 2 ) 3 .3KN0 2 . 
The formation involves oxidation of the cobaltous salt, and 
this is effected by some of the nitrogen trioxide which is 
set free. Thus with the chloride the action may be repre- 
sented as follows : 

CoCl 2 + 7KN0 + 2C 2 H 4 2 = 

2KC1 + Co(N0 2 ) 3 .3KN0 2 + 2KC 2 H 3 2 + H 2 + NO. 

Nickel does not form a similar compound of a nickelic salt, 



640 COLLEGE CHEMISTRY. 

but simply forms a double nitrite, containing the nickelous 
salt Ni(N0 2 ) 2 .4KN0 2 . 

Cobalt compounds color the bead of microcosmic salt 
blue both in the reducing and oxidizing name. Nickel 
colors it reddish brown in the oxidizing flame when hot, 
and pale yellow when cold. In the reducing flame it is 
gray. 

EXPERIMENTS. 
Iron and its Compounds. 

Experiment 239. — Make ferric chloride by heating the purest 
iron wire in a current of chlorine. Also make a solution by dis- 
solving iron in hydrochloric acid, and oxidizing the solution with 
nitric acid. 

Experiment 240. — Make ferrous sulphate by dissolving iron 
in dilute sulphuric acid, and evaporating to crystallization. Dis- 
solve equivalent quantities of ferrous sulphate and ammonium 
sulphate, and evaporate to crystallization. 

Experiment 241. — Dissolve ferric hydroxide in sulphuric acid, 
and evaporate to dryness. 



CHAPTER XXXIV. 

ELEMENTS OF FAMILY VIII, SUB-GROUP B: 
RUTHENIUM.— RHODIUM.— PALLADIUM. 

ELEMENTS OF FAMILY VIII, SUB-GROUP C: 
OSMIUM.— IRIDIUM.— PLATINUM. 

General. — Comparing the members of the three sub- 
groups of Family VIII with reference to their atomic 
weights and specific gravities, we have the following 
remarkable table : 

Fe Co m 

At. Wt. 56 At. Wt. 59 At. Wt. 58.7 

Sp. Gr. 7.8 Sp. Gr. 8.5 Sp. Gr. 8.8 

Ru Rh Pd 

At. Wt. 101.7 At. Wt. 103 At. Wt. 106 
• Sp. Gr. 12.26 Sp. Gr. 12.1 Sp. Gr. 11.5 

Os Ir Pt 

At. Wt. 191 At. Wt. 193 At. Wt. 194.8 

Sp. Gr. 22.48 Sp. Gr. 22.42 Sp. Gr. 21.50 

It will be observed that the atomic weights and specific 
gravities of the members of each sub-group are approxi- 
mately the same. But just as there is a gradual change 
in the chemical conduct from iron to nickel in the iron 
group, so a similar gradation of properties is observed in 
the other two groups. As far as the variety of com- 
pounds which they form is concerned, ruthenium and 
osmium are more like iron than they are like rhodium and 
iridium. Further, rhodium and iridium resemble each 

6di 



642 COLLEGE CHEMISTRY. 

other, as regards the variety of their compounds, more 
closely than they resemble palladium and platinum, and 
a similar resemblance is noticed between palladium and 
platinum. A full discussion of these relationships would 
lead too far. 

The Platinum Metals. 

The six elements of Sub-Groups B and 0, Family 
VIII, are generally grouped together and spoken of as 
the platinum metals. They occur together in nature, and 
almost always in alloys, into the composition of which all 
enter. The chief constituent is platinum, which is present 
to the extent of 50 to 80 per cent, and over. The alloys 
occur in only a few localities, in the Ural Mountains, in 
California, Australia, Borneo, and a few other places, and 
form small pieces which are mixed with sand and earth. 
They generally contain also gold, iron, and copper. Pal- 
ladium occurs, further, in a gold ore which is found in 
Brazil. 

Metallurgy. — The process for obtaining the metals from 
the ores is based mainly upon the following facts: (1) 
Gold is soluble in dilute aqua regia, while platinum 
requires concentrated aqua regia; (2) platinic chloride, 
Pt01 4 , and iridium chloride, IrCl 4 , form, with ammonium 
chloride, difficultly soluble compounds of the formulas 
(NH 4 ) 2 PtCl 6 (PtCl 4 .2NH 4 Cl)and(NH 4 ) 2 IrCl 6 (IrCl 4 .2NH 4 Cl). 
When these compounds are ignited, they are completely 
decomposed, and the metals are left behind. When, 
therefore, platinum-ore has been freed as far as possible 
from sand and earth, it is first treated with dilute aqua 
regia, which removes the gold, and then with concentrated 
aqua regia, which dissolves the platinum together with a 
little iridium, leaving an alloy of iridium and osmium. 

When the solution thus obtained is treated with am- 
monium chloride, both metals are precipitated; and when 
the precipitate is ignited, both metals are left behind in 
the form of a spongy mass, This consists, however, almost 



RUTHENIUM— OSMIUM. 643 

wholly of platinum, the amount of iridium being very 
small. 

Ruthenium, Ru (At. Wt. 101.7). 

Preparation. — Ruthenium is obtained from the residue 
which is left undissolved when platinum-ore is treated 
with concentrated nitro-hydrochloric acid. 

Properties. — When heated in oxygen it burns and forms 
the oxide, Ku0 2 . It is insoluble in the strong acids, and 
even in nitro-hydrochloric acid it is almost insoluble. 
Owing to its power to form salts of ruthenious acid, it is 
dissolved when heated with potassium hydroxide and an 
oxidizing agent, such as saltpetre or potassium chlorate, 
and afterwards treated with water. 



Osmium, Os (At. Wt. 191). 

Preparation. — As stated above, this element is left un- 
dissolved in the form of an alloy with iridium when 
platinum-ore is treated with concentrated nitro-hydro- 
chloric acid. In order to separate it from the iridium, 
advantage is taken of the fact that it forms a volatile 
peroxide, Os0 4 , similar to that formed by ruthenium, 
while iridium does not. 

Properties. — The metal does not melt at the highest 
temperatures reached artificially. It has the highest 
specific gravity of all known substances; is easily oxidized 
when in finely divided condition; and is converted either 
by the oxygen of the air or by nitric acid into osmium 
peroxide, Os0 4 . The metal as well as the oxides forms 
the peroxide, Os0 4 , when heated in the air. This is also 
formed by treating a heated mixture of sodium chloride 
and the alloy of osmium and iridium with chlorine and 
water vapor. It is commonly called osmic acid, though 
its acid properties are very weak. Like ruthenium 
peroxide it is volatile. It sublimes in colorless, lustrous 
needles, and boils without decomposition at a temperature 



644 ' COLLEGE CHEMISTRY. 

a little above 100°. It has an intense odor similar to that 
of chlorine, and its vapor attacks the eyes and respiratory 
organs somewhat as chlorine does. It dissolves slowly in 
water, and reducing agents precipitate the metal from the 
solution. A solution of osmic acid is used in microscopic 
work. When injected into the tissues, the parts are 
hardened and colored. 

Rhodium, Eh (At. Wt. 103). 

Rhodium has no acid properties, and does not form a 
peroxide corresponding to those of ruthenium and osmium. 
On the other hand, its oxide, Rh 2 3 , is basic. The chlo- 
ride RhCl 3 is readily formed. It is doubtful whether the 
di- and tetrachlorides have been made. 

Ikidium, Ir (At. Wt. 193). 

Preparation. — The extraction of iridium with platinum 
and with osmium from platinum-ore was referred to above. 
In order to separate it from platinum, advantage is taken 
of the fact that it forms a trichloride, Ir01 3 , which with 
ammonium chloride gives an easily soluble double chloride. 

Properties. — Iridium has a grayish-white color, and 
resembles polished steel. Its specific gravity is nearly the 
same as that of osmium, being 22.42. It is harder and 
more brittle than platinum; melts at a higher temperature; 
and unless it is finely divided it is not dissolved by nitro- 
hydrochloric acid. When heated with potassium hydroxide 
and saltpetre it is converted into the oxide. 

Palladium, Pd (At. Wt. 106). 

Preparation. — The chief source of palladium is a 
Brazilian gold-ore. From this ore the metal can be 
obtained by various methods, one of which consists in 
melting it together with silver; and then treating it with 
nitric acid, when the silver and palladium dissolve, and 
the gold remains undissolved. The silver is precipitated 



PALLADIUM— PLATINUM. 645 

us chloride and fche palladium as the cyanide, and when 
the latter is ignited it is decomposed, leaving palladium. 

Properties. — Palladium resembles iridium and platinum 
in appearance. Its specific gravity is only about half as 
great as that of platinum, being 11.5; it is more easily 
fusible than platinum, and dissolves in nitric acid and in 
hot concentrated sulphuric acid. The property of pal- 
ladium which has perhaps attracted most attention is its 
power to absorb hydrogen, and form 

Palladium-Hydrogen.— The formation of this compound 
was referred to under Hydrogen (wiiich see). The com- 
bination takes place even at the ordinary temperature, but 
best at 100°. If the metal is brought into hydrogen at 
this temperature, it absorbs more than 900 times its 
volume, forming an alloy of the composition Pd 2 H. This 
alloy has a greater volume and lower specific gravity than 
the palladium from which it is formed. At 130° it begins 
to decompose under the atmospheric pressure, but con- 
tinued heating at a red heat is necessary to decompose it 
completely. If allowed to lie in contact with the air the 
hydrogen is oxidized to water. Palladium-hydrogen acts 
as a strong reducing agent, the hydrogen which it gives 
up being apparently in the nascent or atomic condition. 

Platinum, Pt (At. Wt. 194.8). 

Preparation. — A general idea of the method of pro- 
cedure in extracting platinum from its ores was given on 
p. 642. Thus prepared, however, it always contains 
iridium, and for some purposes for which platinum is used 
this is objectionable. In order to purify the metal advan- 
tage is taken of the fact that iridium chloride can be con- 
verted into a trichloride, which with ammonium chloride 
forms an easily soluble double salt. The metal as ob- 
tained by igniting ammonium platinic chloride forms a 
gray spongy mass known as spongy platinum. When 
a solution of platinous chloride is boiled with potassium 
hydroxide, and alcohol gradually added, the salt is re- 



646 COLLEGE CHEMISTRY. 

duced, and the platinum is precipitated as an extremely 
fine powder, known as platinum Mack. When spongy 
platinum and platinum-black are heated to fusion by the 
oxyhydrogen flame they are converted into the compact 
variety. 

Properties. — Platinum is a grayish-white metal resem- 
bling polished steel; it can be drawn out into very fine 
wire; it melts in the flame of the oxyhydrogen blowpipe, 
and when heated above its melting-point it is volatile: its 
specific gravity is 21.5. At white heat it can be welded. 
It is not dissolved by nitric acid, hydrochloric acid, or 
sulphuric acid, but it dissolves in nitro-hydrochloric acid, 
forming the acid, H 2 PtCl 6 . Fusing akalies, and particu- 
larly a mixture of caustic potash and saltpetre, act upon 
it; but the alkaline carbonates do not. In contact with 
red-hot charcoal and silicon dioxide a compound of silicon 
and platinum is formed. Finely divided platinum has to 
a remarkable extent the power of condensing gases upon 
its surface. It absorbs, for example, 200 times its own 
volume of oxygen, and other gases in a similar way. The 
oxygen thus absorbed is in an active condition, and if 
oxidizable substances are brought in contact with it they 
are easily oxidized. Thus when a current of hydrogen is 
allowed to flow against a piece of spongy platinum it takes 
fire, owing to the presence of the condensed oxygen in the 
pores of the platinum. Similarly, when sulphur dioxide 
and oxygen are allowed to flow together over spongy 
platinum, or even the compact metal, the two gases unite 
to form sulphur trioxide. 

Applications of Platinum. — The metal is of great value 
to the chemist on account of its power to resist high tem- 
peratures and the action of most chemical substances. It 
is used in the laboratory in the form of wire, foil, cruci- 
bles, evaporating-dishes, tubes, etc., etc. From what was 
said above it cannot be used with alkalies and saltpetre, 
nor with nitro-hydrochloric acid. Platinum vessels, 
further, should not be placed upon red-hot charcoal. 
Metallic salts which are easily reduced, such as those of 



ALLOYS OF PLATINUM. 647 

antimony and bismuth, should not bo heated in platinum 
vessels, as the reduced elements, like silicon, form alloys 
with the platinum, and these, as a rule, are easily fusible. 
In the concentration of sulphuric acid on the large scale 
platinum stills are used. The price of platinum is not as 
high as that of gold, but much higher than that of silver. 

Alloys of Platinum. — The only alloy of platinum which 
is of any special importance is that which it forms with 
iridium. A small percentage of iridium diminishes the 
malleability of platinum very markedly, and makes it 
brittle; it, however, increases its resistance to the action 
of reagents. An alloy of 90 per cent platinum and 10 per 
cent iridium has been adopted by the French Government 
as the best material from which to make normal meters. 
This alloy is very hard, as elastic as steel, more difficultly 
fusible than platinum, entirely unchangeable in the air, 
and is capable of a high polish. 

Chlorides. — Like palladium, platinum forms two chlo- 
rides, platinous chloride, PtCl 2 , and jy/a^mc cldoride, 
PtCl 4 . Platinous chloride is formed by passing chlorine 
over platinum sponge heated to 240° to 250°. Platinic 
chloride is formed by heating chlorplatinic acid, H 2 PtCl 6 , 
in a current of chlorine to 360°. Platinic chloride is 
easily soluble in water. 

Chlorplatinic Acid, H 2 PtCl 6 , is formed by direct union 
of platinic chloride with hydrochloric acid, and by dissolv- 
ing platinum in aqua regia and evaporating the solution. 
It crystallizes with six molecules of water, and forms a 
series of salts called the clilorplatinates, to which reference 
has already been made. Those most commonly met with 
in the laboratory are the potassium salt, K 2 PtCl 6 , or 
PtCl 4 .2KCl, and the ammonium salt, (NHJ 3 PtCl 6 , or 
Pt01 4 .2NH 4 01, both of which are difficultly soluble in 
water, and are therefore precipitated when chlorplatinic 
acid (platinic chloride) is added to solutions containing 
potassium or ammonium chloride. The sodium salt is 
easily soluble in water. Chlorplatinic acid appears to be 
analogous to fluosilicic acid, H SiF . 



648 COLLEGE CHEMISTRY. 

Sulphides. — There are two sulphides of platinum which 
are analogous to the two oxides, PtO and Pt0 2 . These 
are plalinous sulphide, PtS, and platinie sulphide, PtS 2 . 
They are black insoluble compounds, which are precipi- 
tated when hydrogen sulphide or soluble sulphides are 
added to solutions of platinous and platinic chlorides. 

EXPERIMENT. 

Platinum. 

Experiment 242. — Prepare a solution of chlorplatinic acid as 
follows : Heat platinum in a flask with concentrated nitric acid, 
adding from time to time a few drops of hydrochloric acid. 
After the metal is dissolved evaporate to dryness on a water-bath. 
Dissolve in water and filter. 

Chemical Analysis. 

At the end of each chapter treating of the metallic or base- 
forming elements there is given a list of such reactions as are of 
special value for analytical purposes, together with such explana- 
tory statements as seem called for. The student who is engaged 
in analytical work will generally find these explanations sufficient 
to enable him to keep his ideas clear in regard to the reactions 
with which he is dealing, provided, at the same time, he carefully 
studies the chapter to which the explanations form an appendix. 
As an introduction to analytical work, a general study of chem- 
ical reactions is necessary, and the fuller this is the better. There 
are many small books in existence in which good directions are 
given for work of this kind. The student is, however, advised 
to supplement the book he may be using by such experiments as 
may suggest themselves on reading the corresponding chapters 
in this book. The directions there found will generally be quite 
sufficient for the purpose, and it is therefore not considered nec- 
essary to give more specific directions in this place. In general, 
the more the student ocsupies himself in the laboratory with 
chemical substances the more rapidly will his chemical ideas 
grow. But it is necessary that he should avoid working by " rule 
of thumb," and for this purpose constant reference to some larger 
text-book in which the relations between the substances and the 
reactions he is dealing with are discussed in a broad way is of the 
highest importance. 



CHAPTER XXXV. 
SOME FAMILIAR COMPOUNDS OF CARBON. 

Organic Chemistry. — When the compounds that are 
obtained from plants and animals were first studied, they 
were supposed to be entirely different from the compounds 
obtained from the inorganic, or mineral, constituents of 
the earth. The former were called organic compounds 
because they were obtained from organized things; while 
the latter were called inorganic compounds. Organic 
compounds were the subject of Organic Chemistry, and 
inorganic compounds formed the subject of Inorganic 
Chemistry. These names are still in use, though they 
have lost their original meaning. Organic Chemistry 
now means only the Chemistry of the Compounds of Carbon. 

Occurrence of the Compounds of Carbon. — The com- 
pound of carbon that occurs most widely distributed in 
nature is carbon dioxide. This, as has been pointed out, 
is the starting-point of all life on the globe. All living 
things are formed from it either directly or indirectly. 
Attention has been called to the fact that starch and 
cellulose are the principal compounds found in plants, and 
that fats, albumin, and fibrin are the most common sub- 
stances found in animals. 

Formation of Hydrocarbons. — Certain natural processes 
which are not thoroughly understood have given rise to 
the formation of a complex mixture of organic com- 
pounds, principally hydrocarbons, in petroleum. 

Distillation of Coal. — The destructive distillation of coal 
for the purpose of making illuminating-gas, and the 
formation of coal-tar, have already been referred to. 

649 



6$° COLLEGE CHEMISTRY. 

Coal-tar is one of the most important sources of compounds 
of carbon. The hydrocarbons benzene, C 6 H 6 , toluene, 
C 7 H 8 , xylene, C 8 H 10 , naphthalene, C 10 H 8 , anthracene, 
C U H 10 , etc., are obtained from this source. 

Distillation of Wood. — Wood is heated in closed vessels 
mostly for the purpose of making charcoal, as already 
explained. Among the products obtained from this source 
are wood-spirit, or methyl alcohol, and pyroligneous acid, 
or acetic acid. Large quantities of acetic acid are prepared 
in this way. 

Distillation of Bones. — In order to make bone-black, 
bones are subjected to destructive distillation. The oil 
which passes over is collected and known as bone-oil. 
This is the source of a large number of compounds which 
are of special interest on account of their connection with 
the valuable alkaloids quinine, morphine, etc. 

Fermentation. — A number of the most important com 
pounds of carbon are formed by a process known as fermen- 
tation. This is a general term meaning any process in 
which a chemical change is effected by means of minute 
animal or vegetable organisms. The best -known example 
of fermentation is that of sugar, which gives rise to the 
formation of ordinary alcohol. 

Classes of Compounds of Carbon. — The chief classes of 
these compounds are the hydrocarbons, the alcohols, the 
aldehydes, the acids, the ethers, and the ethereal salts. 
First a few of the best-known examples of each of these 
classes will be taken up, and afterwards some other familiar 
compounds which do not belong to any one of these classes. 

HYDKOCAKBOKS. 

Compounds of Carbon and Hydrogen. — It is not an easy 
matter to effect combination between carbon and hydrogen 
in the laboratory except in a few simple cases. In nature 
processes are in operation which give rise to the formation 
of a large number of compounds containing these elements; 
and, further, in the manufacture of illuminating-gas from 



PETROLEUM. 651 

Coal the conditions are such as fco cause the combination 
of carbon and hydrogen, several interesting compounds 
being thus formed. There are no other two elements 
which combine with each other in as many different pro- 
portions as carbon and hydrogen. The compounds thus 
formed are known as hydrocarbons. The number of 
hydrocarbons known is large, being somewhere near two 
hundred. Fortunately, investigation has shown that quite 
simple relations exist between these compounds; and 
hence, though the number is large, the study is not as 
difficult as might be expected. 

Petroleum is an oily liquid found in many places in the 
earth in large quantity, particularly in Pennsylvania and 
the Caucasus. In the earth it contains both gases and 
liquids. When it is brought into the air, the pressure 
being removed, the gases are given off. There are several 
gaseous hydrocarbons given off, and a large number of 
liquids left behind. 

Refining of Petroleum.— The vapors from petroleum 
when mixed with air are explosive, and the thicker liquids 
clog the lamps and wicks. Therefore these must be 
removed before the oil is fit for household use. This is 
done by (1) distilling, (2) washing with sulphuric acid, 
(3) washing with alkali, and (4) washing with water. The 
product thus prepared is called kerosene. 

In refining petroleum a number of products are obtained 
which cannot be used in lamps. Those which are lighter 
than kerosene, that is to say those which boil at a lower 
temperature, are known as gasoline, naphtha, benzine, etc. 
From the heavier portions, or those which boil at higher 
temperatures than kerosene, lubricating oils and paraffin 
are made. Each of these substances is a mixture of 
several chemical compounds. 

Hydrocarbons contained in Petroleum. — The simplest 
hydrocarbon contained in petroleum is methane, or marsh- 
gas, CH 4 ; the next has the composition C 2 H 6 , the next 
C 3 H 8 , etc. It will be seen that these compounds bear a 
simple relation to one another, as far as composition is 



^52 COLLEGE CHEMISTRY. 

concerned. They are the first members of a series the 
names and symbols of the first eight members of which are 
given below: 

CH 4 , Methane, or Marsh-gas; 

C 2 H 6 , Ethane; 

C 3 H 8 , Propane ; 

C 4 H 10 , Butane; 

C 5 H 12 , Pentane; 

C 6 H U , Hexane; 

C 7 H 16 , Heptane; 

C 8 H 18 , Octane. 

Homology. — The first member of the series differs from 
the second by CH 2 ; there is also this same difference, in 
general, between any two consecutive members of the 
series. This relation is known as homology, and such a 
series as an homologous series. Carbon is distinguished 
from all other elements by its power to form homologous 
series. 

The Ethylene Series of Hydrocarbons, — Besides the 
series above mentioned, which is known as the marsh-gas 
series, there are other homologous series of hydrocarbons. 
There is one beginning with ethylene, C 2 H 4 , examples of 
which are 

Ethylene, C 2 H 4 ; 

Propylene, C 3 H 6 ; 

Butylene, C 4 H 8 . 

The Acetylene Series, — There is a series beginning with 
acetylene, examples of which are 

Acetylene, C„H 2 ; 
Allylene, C 3 H/ 

The Benzene Series. — Another series begins with ben- 
zene, C JL. Some of the members of this series are 

' DO 

Benzene, C 6 H 6 ; 
Toluene. C 7 H 8 ; 
Xylene, 8 H 10 . 



MARSH-GAS. 653 

Marsh-gas, Methane, Fire-damp, CH 4 .— Marsh-gas is 
found in nature in petroleum, and is given off when the 
oil is taken out of the earth, and the pressure is removed. 
It is formed, as the name implies, in marshes, as the 
product of a reducing process. Vegetable matter is com- 
posed of carbon, hydrogen, and oxygen. When it under- 
goes decomposition in the air in a free supply of oxygen. 
the final products formed are carbon dioxide and water. 
AVlien the decomposition takes place without access of 
oxygen, as under water, marsh-gas, which is a reduction- 
product, is formed. The gas can be made in the laboratory 
by treating aluminium carbide, C 3 A1 4 , with water: 

C,A1, -j- 12H 2 = 3CH 4 + 4A1(0H) 3 . 

It can also be made by passing a mixture of hydrogen sul- 
phide and the vapor of carbon bisulphide over heated 
copper. 

Marsh-gas is met with in coal-mines, and is known to 
the miners as fire-damp, "damp" being the general 
name applied to a gas, and the name fire-damp meaning a 
gas that burns. To prepare it in the laboratory, it is 
most convenient to heat a mixture of sodium acetate and 
quicklime. The change which takes place will be readily 
understood by considering it as a simple decomposition of 
acetic acid. Acetic acid has the formula C 2 H 4 2 . When 
heated alone, it boils and does not suffer decomposition. 
If it is converted into a salt, and heated in the presence 
of a base, it breaks down into marsh T gas and carbon 
dioxide : 

C 2 H 4 2 = CH 4 + CO,. 

The carbon dioxide, which forms salts with bases, does not 
pass off, but remains behind in the form of a salt of car- 
bonic acid. 

Marsh-gas is a colorless, transparent, tasteless, inodorous 
gas. It is slightly soluble in water. It burns, forming 
carbon dioxide and water. When mixed with air, the 



654 COLLEGE CHEMISTRY. 

mixture explodes if a flame or spark comes in contact with 
it. This is one of the causes of the explosions which so 
frequently occur in coal-mines. To prevent these explo- 
sions a special lamp was invented by Sir Humphry Davy, 
which is known as Davy's safety-lamp (p. 374). 

Substitution-products of the Hydrocarbons. — Marsh -gas 
and other hydrocarbons undergo change when treated with 
chlorine and bromine. The change consists in the substi- 
tution of one or more atoms of chlorine or of bromine for 
the same number of atoms of hydrogen. In the case of 
marsh-gas and chlorine the possible changes are repre- 
sented as below: 

CH 4 + 01, = CH 3 C1 + HOI 
CH 3 01 + Cl 2 = CH 2 01 2 + HC1 
CH 2 C1 2 + 01 2 = CHC1 3 + HC1 
0HC1 3 + Cl 2 = CC1 4 + HOI. 

All the products represented are known. 

Chloroform, CHC1 3 . — Chloroform can be made as above 
indicated, but it is made on the large scale by treating 
alcohol (which see) or acetone (which see) with bleaching- 
powder. It is a heavy liquid with an ethereal odor and a 
somewhat sweet taste. It is one of the most valuable 
anaesthetics, though there is some danger attending its use. 

Iodoform, CHI 3 . — This compound, like chloroform, is a 
substitution-product of marsh-gas. It is made by bring- 
ing together alcohol, an alkali, and iodine. It is a solid 
substance, soluble in alcohol and ether, but insoluble in 
water. It crystallizes in six-sided yellow plates. It is 
extensively used as a dressing for wounds in surgery. 

Ethylene, Olefiant Gas, 2 H 4 . — This hydrocarbon is 
formed by heating a mixture of ordinary alcohol and con- 
centrated sulphuric acid. The reaction is represented 
thus: 

C 2 H 6 = H 2 + C,H 4 . 

Alcohol. Ethylene. 

Ethylene is a colorless gas, which can be condensed to 



ALCOHOLS. 655 

a liquid. It burns with a luminous flame. With oxygen 
it forms an explosive mixture. 

Acetylene, C 2 H 2 . — Acetylene is formed when a current 
of hydrogen is passed between carbon poles, which are in- 
candescent in consequence of the passage of a powerful 
electric current. In this case carbon and hydrogen com- 
bine directly. It is formed also when the flame of an 
ordinary laboratory gas-burner (Bunsen burner) "strikes 
back/' or burns at the base without a free supply of air. 
It is now made on the large scale by treating calcium car- 
bide, CaC 2 , with water: 

CaC 2 + H 2 = C 2 H 2 + CaO. 

Its odor is unpleasant. It burns with a luminous, 
smoky flame. 

ALCOHOLS. 

Methyl Alcohol, Wood-spirit, CH 4 0. — This is formed in 
the distillation of wood together with other products. It 
has, when pure, a pleasant odor and taste, and acts upon 
the animal system very much as ordinary alcohol does. It 
burns without giving light or smoke, and can therefore 
be used in lamps for heating-purposes as ordinary alcohol 
is. It is used in the manufacture of varnishes. 

Ethyl Alcohol, Spirits of Wine, 2 H 6 0.— This well- 
known substance is formed by the fermentation of grape- 
sugar or glucose. 

What Change takes Place in the Sugar ? — If the solu- 
tion in the flask is examined carefully it will be found to 
contain alcohol and no sugar. Grape-sugar has the com- 
position expressed by the formula C 6 H 12 6 . By fermenta- 
tion it is decomposed, forming alcohol, C 2 H 6 0, and carbon 
dioxide, C0 2 , thus: 

C 6 H 12 6 = 2C 2 H 6 + 2C0 2 . 

What Causes the Change ? — It has been found that the 
change of grape-sugar is caused by small organized bodies 



656 COLLEGE CHEMISTRY. 

which grow in the solution. These bodies are contained 
in ordinary yeast. 

Germs in the Air. — When fruit-juices that contain sugar 
are exposed to the air they undergo fermentation with- 
out the addition of yeast. This is due to the fact that 
the germs or seeds of the bodies which cause fermentation 
are everywhere floating in the air. Hence when a liquid in 
which these seeds can grow is exposed to the air, the 
bodies are formed and fermentation takes place. 

Different Kinds of Fermentation. — The fermentation 
which yields alcohol is only one of many kinds. Among 
the others are: (1) lactic-acid fermentation, which takes 
place in the souring of milk; and (2) acetic-acid fermenta- 
tion, which causes the transformation of alcohol into acetic 
acid. The latter ferment is contained in " mother of 
vinegar." 

Distillation of Fermented Liquids, — In order to get the 
alcohol from liquids which have undergone fermentation 
they must be distilled. For this purpose very perfect 
forms of stills have been devised, so that the alcohol passes 
over nearly free from other substances. Usually it con- 
tains impurities known as fusel oil. 

Properties of Alcohol. — Pure ethyl alcohol has a peculiar, 
pleasant odor. It remains liquid at very low temperatures, 
but has been converted into a solid at a temperature of 
— 130.5°. It burns with a flame which does not deposit 
soot, and was hence formerly much used in laboratories 
for heating purposes, and is still used where gas cannot be 
obtained. Its effects upon the human system are well 
known. It intoxicates when taken in dilute form, while 
in large doses it is poisonous. It lowers the temperature 
of the body when taken internally, although it causes a 
sensation of warmth. 

Uses of Alcohol. — Alcohol is the principal solvent for 
organic substances. It is hence extensively used in the 
arts, as in the manufacture of varnishes, perfumes, and 
tinctures of drugs. Most beverages in use owe their in- 
toxicating-power to the presence of alcohol. The milder 



GL YCERIN- ALDEH YDES. 657 

forms of beer contain from 2 to 3 per cent; light wines 
about 8 per cent; while whiskey, brandy, etc., sometimes 
contain as much as GO to 75 per cent. 

Glycerin, C 3 H 8 3 . — Glycerin is an alcohol which occurs 
very widely distributed as a constituent of fats. The 
relation it bears to the fats will be explained when the 
acids which enter into the fats are taken up. It is obtained 
from the fats by boiling them with an alkali like caustic 
soda or caustic potash, or by heating with steam. 

Properties. — Glycerin is a thick, colorless liquid with a 
sweetish taste. It attracts moisture from the air, and is 
hence used to keep surfaces moist. 

ALDEHYDES. 

Acetic Aldehyde, Ordinary Aldehyde, C 2 H 4 0. — This 
compound is formed by oxidizing ordinary alcohol, the 
change being represented by this equation : 

C 2 H 6 + = C 2 H 4 + H 2 0. 

Aldehyde is a volatile liquid with a characteristic pene- 
trating odor. When left to itself, and especially when 
treated with a number of other things, it is converted into 
another substance of the same composition. This is called 
paraldehyde. A determination of the molecular weight of 
the substance by the method of Avogadro has shown that 
it must be represented by the formula C 6 H 12 3 . The 
change from aldehyde to paraldehyde must, therefore, be 
represented thus : 

3C 2 H ( = C 6 H 12 3 . 

Paraldehyde is used in medicine. 

Chloral, C 2 C1 3 H0, is a compound formed by the action 
of chlorine on alcohol. It is related to aldehyde, as 
chloroform is related to marsh-gas, that is to say, it is a 
trichlorine substitution-product. It is a colorless liquid. 
With water it forms a crystallized confound, chloral 
hydrate, C 2 Cl ? HO -f- H 2 0, which is easily soluble in water, 
and crystallizes from the solution in colorless prisms. 



658 COLLEGE CHEMISTRY. 

Taken internally in doses of from 1.5 to 5 grams, it 
produces sleep. In larger doses it acts as an anaesthetic. 

ACIDS. 

Formic Acid, CH 2 2 . — This acid occurs in nature in red 
ants, in stinging-nettles, in the shoots of some of the 
varieties of pine, and elsewhere. It is a colorless liquid. 
Dropped on the skin, it causes extreme pain and produces 
blisters. 

Acetic Acid, C 2 H 4 2 . — This is the acid contained in 
vinegar, and the value of vinegar is due to its presence. 
It is formed from alcoholic liquids by exposing them to the 
air, in consequence of the presence of a microscopic 
organism which is contained in what is commonly known 
as " mother of vinegar." The formation of acetic acid 
from alcohol is due to the action of oxygen as represented 
in the equation 

C 2 H 6 + 2 = C 2 H 4 2 + H 2 0. 

Alcohol. Acetic acid. 

But oxygen alone does not effect the change. "When the 
ferment is present the oxidation takes place. Acetic acid 
is also obtained by distilling wood. Hence the names 
pyroligneous acid and wood-vinegar. 

Properties. — Acetic acid is a clear, colorless liquid. It 
has a very penetrating, pleasant, acid odor, and a sharp 
taste. The pure substance acts upon the skin like formic 
acid, causing pain and raising blisters. 

Uses. — Acetic acid is extensively used, chiefly in the 
dilute form known as vinegar. It is used in calico-print- 
ing in the form of iron and aluminium salts. With iron 
it gives hydrogen, which is needed in the manufacture of 
certain compounds used in making dyes. 

Salts of Acetic Acid. — The best-known salts of acetic 
acid are lead acetate, Pb(0 2 H 3 2 ) 2 , commonly called sugar 
of lead; and copper acetate, Cu(C 2 H 3 2 ) 2 , a variety of 
which is known as verdigris. 

Fatty Acids. — Formic and acetic acids are the first 
members of an homologous series (see p. 652). Some of 



FATTY ACIDS. 659 

the more important members are named in the following 
table : 

Formic acid CH a O a . 

Acetic " C a H 4 O a . 

Propionic " C 3 lI 6 O a . 

Butyric " C 4 H 8 O a . 

Palmitic " C 16 H 32 02. 

Stearic " C ie H 36 O a . 

They are called fatty acids for the reason that many of 
them are obtained from fats. 

Butyric acid, C 4 H 8 2 , is of special interest because it is 
obtained from butter by boiling with caustic potash. It 
occurs also in many other fats. There is a butyric-acid 
ferment contained in putrid cheese which has the power 
of converting sugar into butyric acid. 

Palmitic acid, C 16 H 32 2 , is obtained from many fats, but 
palm-oil is especially rich in it. 

Stearic acid, C 18 H 36 2 , is the acid contained in the fat 
kuown as stearin. The so-called "stearin candles" are 
made of a mixture of palmitic and stearic acids. 

Soaps. — Soaps are the alkali salts of the acids contained 
in fats, especially of palmitic and stearic acids. Fats are 
compounds of these acids with glycerin. When the fats 
are boiled with an alkali, as caustic soda, the correspond- 
ing salts of the acids are formed, while the glycerin is set 
free. The potassium and sodium salts of palmitic and 
stearic acids are the soaps. 

Use of Soap. — The cleansing power of soap depends upon 
the fact that it dissolves the oily film on the surface of the 
skin and thus facilitates the removal of the foreign sub- 
stances commonly known as dirt. 

Action of Soap on Hard Waters. — As has been explained, 
a hard water is one that contains salts in solution. Tem- 
porary hardness is that which is caused by calcium car- 
bonate held in solution in the water by carbon dioxide. 
Permanent hardness is caused by calcium sulphate or 
magnesium salts. The calcium and magnesium salts of 
palmitic and stearic acids are insoluble in water. There- 



660 COLLEGE CHEMISTRY. 

fore, when soap is added to a hard water these insoluble 
salts are precipitated and give the water a hard feeling. 
In attempting to wash the hands with soap in a hard water 
they become covered with a thin layer of the insoluble 
salts which prevents them from rubbing freely over each 
other, and makes them feel sticky. Before the soap can 
do any good all the lime-salt must be precipitated. The 
action in the case of temporary hardness is represented by 
the equation 

2Na0 16 H 31 2 + CaCO, = Ca(C 16 H 31 2 ) 2 + Na 2 C0 3 . 

Soap. Calcium palmitate. 

In the case of permanent hardness it is represented by 
the equation 

2NaC 16 H 31 2 + CaS0 4 = 0a(0 16 H 31 O 2 ) 2 + Na 2 S0 4 . 

Relations of the Soap Industry to other Industries. — A 

great chemist and philosopher has said that the quantity 
of soap used in a country is a measure of the civilization 
of that country. Certain it is that soap is only used by 
civilized people, and that by them it is used in enormous 
quantities. In many farm-houses a primitive method for 
the manufacture of soap is practised, consisting in treating 
refuse fats with the lye extracted from wood-ashes. A 
soft soapy mass is thus obtained known as " soft-soap." 
Fats form the starting-point in the manufacture of all soap. 
These are generally treated with caustic soda. Caustic soda 
is all made from sodium carbonate by the action of lime; 
and, as has been seen, sodium carbonate is made from 
common salt mostly by the Le Blanc process, which re- 
quires sulphuric acid. Thus the manufacture of sulphuric 
acid and sodium carbonate is intimately related to the 
manufacture of soap. 

Oxalic Acid, C 2 H 2 4 .— This acid occurs very widely dis- 
tributed in nature, as in the sorrels, which owe their acid 
taste to the presence of acid potassium oxalate, KC 2 H0 4 ; 
and as the ammonium salt in guano. It is probably one 
of the first substances formed from carbon dioxide in the 



ETHERS. 66 1 

plant. It is manufactured by heating wood shavings or 
sawdust with caustic soda and caustic potash. Oxalic acid 
is an active poison. It is used in calico-printing, and in 
cleaning brass and copper surfaces. 

Lactic Acid, C 3 II 6 3 . — Lactic acid is made by the fer- 
mentation of sugar by means of the lactic-acid ferment. 
The reaction effected by the ferment is represented by the 
equation 

C 6 H I2 6 = 2C 3 H 6 3 . 

Malic Acid, C 4 H 6 5 . — This acid is very widely distributed 
in the vegetable kingdom, as in apples, cherries, etc. 

Tartaric Acid, C 4 H 6 6 . — Tartaric acid occurs very widely 
distributed in fruits, sometimes uncombined, sometimes in 
the form of the potassium or calcium salt; as, for example, 
in grapes, berries of the mountain-ash, potatoes, cucum- 
bers, etc., etc. It is prepared from " cream of tartar." 
This is acid potassium tartrate, which is formed when 
grape-juice ferments. 

Citric Acid, C 6 H 8 7 . — Citric acid, like malic and tartaric 
acids, is very widely distributed in nature in many varieties 
of fruit, especially in lemons. It is also found in currants, 
whortleberries, raspberries, gooseberries, etc., etc. It is 
prepared from lemon-juice: 100 parts of lemons yield 5^ 
parts of the acid. It is a solid, crystallized substance, 
soluble in water. It is frequently used for the purpose of 
making lemonade without lemons, and there is no objec- 
tion to its use for this purpose. 

ETHEKS. 

Ether, C 4 H 10 O. — Ordinary ether is the best-known repre- 
sentative of the class of compounds called ethers. It is 
formed from ordinary alcohol by treating it with sulphuric 
acid and distilling. The result of the action which takes 
place is represented by the equation 

2C 2 H 6 = C 4 H 10 O + H 2 0. 

Alcohol. Ether. 



662 



COLLEGE CHEMISTRY. 



Ether is a liquid which boils at a low temperature and 
takes fire and burns readily. Inhaled it produces insensi- 
bility to pain. It is therefore called an ancesthetic. 

• ETHEEEAL SALTS. 

Action of Acids upon Alcohols. — When an acid acts upon 
an alcohol it is neutralized, though not as readily as when 
it acts upon a base. The product is a substance which 
resembles a salt and is called an ethereal salt. Thus when 
nitric acid acts upon alcohol this reaction takes place: 

C 2 H 6 + HN0 3 = C 2 H 5 N0 3 + H 2 0. 

The product C 2 H 5 N0 3 , called ethyl nitrate, is an ethereal 
salt. The alcohol acts as if it were a substance like caustic 
potash and made up thus: C 2 H 5 OH. The resemblance 
between its action and that of caustic potash is shown by 
the equations 

KOH + HN0 3 = KN0 3 + H 2 0, and 

C 2 H 5 OH + HN0 3 = C 2 H 5 N0 3 + H 2 0. 

Saponification. — When an ethereal salt is boiled with a 
caustic alkali it is decomposed, the products being an 
alcohol and an alkali salt. Thus when ethyl nitrate is 
boiled with caustic potash, potassium nitrate and alcohol 
are formed : 

C 2 H 5 N0 3 + KOH = C 2 H 5 OH + KN0 3 . 

This process is called saponification, because the most 
important example is furnished by soap-making. 

Fats. — The fats are ethereal salts in the formation of 
which glycerin, as the alcohol, and three acids take part. 
The three acids are palmitic and stearic acids, already 
mentioned, and oleic acid, C 18 H 34 2 . Although the com- 
position of these substances is comparatively complex, the 
way they act upon one another is simple, and is the same 
as the action of nitric acid upon alcohol in forming ethyl 
nitrate. The fats, then, are the palmitate, stearate, and 
oleate of glyceryl, which bears to glycerin very much the 



ETHEREAL SALTS, ETC. 663 

same relation that ethyl, C 2 H 5 , bears to alcohol. When a 
fat is boiled with caustic soda, glycerin and the sodium 
salts of the acids contained in the fat are formed. 

Butter consists of ethereal salts of glycerin and several 
fatty acids, among which are palmitic, stearic, and butyric. 
Oleomargarin is an artificial butter made from other fats 
than that of milk. 

Ethereal Salts as Essences. — The ethereal salts generally 
have pleasant odors, and it is to their presence that many 
fruits owe their flavors. Some of the compounds are now 
made artificially and used instead of the fruit-extracts. 
Thus the ethyl salt of butyric acid is used under the name 
of essence of pineapples, and the amy I salt of valeric acid 
under the name of essence of apples. 

Nitroglycerin. — Among the more important ethereal 
salts of glycerin are the nitrates. Two of these are 

ro.N0 2 

known, viz., the mono-nitrate, C H. \ OH , and the tri- 

D (OH 

nitrate, C 3 H 5 (0.1Sr0 2 ) 3 , the latter being the chief con- 
stituent of nitroglycerin. Nitroglycerin is prepared by 
treating glycerin with a mixture of concentrated sulphuric 
and nitric acids. It is a pale yellow oil which is insoluble 
in water. At — 20° it crystallizes in needles. It explodes 
very violently by concussion. It may be burned in an 
open vessel, but if heated above 250° it explodes. 
Dynamite is infusorial earth* impregnated with nitro- 
glycerin. Nitroglycerin is the active constituent of a 
number of explosives. 

RELATIONS BETWEEN THE COMPOUNDS CONSIDERED. 

Comparison of the Formulas. — On comparing the formu- 
las of the hydrocarbons of the marsh-gas series (see p. 652) 
with those of the simplest alcohols and the fatty acids, it 
will be seen that these compounds are all related in a 

* That is to say, earth made up of the microscopic flinty shells 
which constitute the fossil remains of certain minute and simple 
plants. 



664 COLLEGE CHEMISTRY. 

simple way. Below are tlie formulas of a few of the 
hydrocarbons, alcohols, and acids: 



drocarbons. 


Alcohols. 


Acids. 


CH 4 


CH 4 


cha 


C 2 H 6 


C 2 H e O 


<W>, 


C S H 8 


S H 8 


C 3 H 6 2 


C 4 H 10 , etc. 


C 4 H 10 O, etc. 


C,H 8 2) 



etc. 



Each of these series is an homologous series. 

Alcohols. — Alcohols have been shown to be derived from 
the hydrocarbons by the replacement of one or more 
hydrogen atoms by oxygen and hydrogen, OH, called 
hydroxyl, or from water by replacing one of the hydrogen 
atoms of the water by a compound of carbon and hydro- 
gen. An alcohol, then, is a hydroxide, just as a metallic 
base is; only, instead of consisting of a metal in combina- 
tion with hydrogen and oxygen, it consists of a compound 
of carbon and hydrogen in combination with hydrogen and 
oxygen. Thus : 

Metallic Bases. Alcohols. 

K(OH) CH 3 (OH) 

Na(OH) C 2 H 5 (OH) 

More Complex Alcohols. — Just as lime is a more complex 
base than caustic potash, as shown by the formulas KOH 
and Oa0 2 H 2 or Ca(OH) 2 , so there are more complex 
alcohols than ordinary alcohol. A good example is fur- 
nished by glycerin, 3 H 8 3 , which has been shown to be 
a hydroxide corresponding to aluminium hydroxide, 
Al(OH) 3 , a fact which is represented by the formula 
C 3 H 5 (OH) 3 . It may be called glyceryl hydroxide, the 
complex, C 3 H 5 , being known as glyceryl. 

Radicals or Residues. — The compounds of hydrogen and 
carbon contained in the alcohols are called radicals or 
residues. We may say that an alcohol is water in which 
a radical has been substituted for half of the hydrogen. 



ACIDS. 



665 



HOH 


C 2 H 5 OH 


Water. 


Ordinary alcohol. 


HOH 


( OH 


HOH 


C3IIJ OH = C.ILO 


HOH 


( OH 


Water. 


Glycerin. 



Acids. — Just as the alcohols have been shown to be 
derived from water, so the organic acids have been 
shown to be derived from carbonic acid. Carbonic 
acid itself is not known. But the carbonates are de- 

rived from an acid of the formula H 2 C0 3 , or CO ■! ^tt. 

If, in this acid, a hydroxyl is replaced by a radical, as, for 
example, by ethyl, C 2 H 5 , a substance of the formula 

CO j qA or C 3 H 6 2 is the result. If methyl, CH 3 , is 

inn- 
OH 3 or 
C S H 4 2 , which is acetic acid. In a similar way all the 
organic acids are derived from carbonic acid. 



EXPERIMENTS. 

Fermentation. 

Experiment 243. — Dissolve about 150 grams of commercial 
grape-sugar in li litres of water in a flask. Connect the flask by 




means of a bent glass tube with a cylinder or bottle containing 
clear lime-water. The vessel containing the lime-water must be 
provided with a cork with two holes. Through one of these 



666 COLLEGE CHEMISTRY. 

passes the tube from the fermentation-flask ; through the other 
a tube connected with a vessel containing solid caustic potash, 
the object of which is to prevent the air from acting upon the 
lime-water. The arrangement of the apparatus is shown in 
Fig. 79. Now add to the solution of grape-sugar a little fresh 
brewers' yeast ; close the connections and allow to stand. Soon 
an evolution of gas will begin, and, as this passes through the 
lime-water, a precipitate will be formed which can be shown to 
be calcium carbonate. 

Aldehyde. 

Experiment 244. — In a small flask put a few pieces of potassium 
bichromate, K2O2O7 > and pour upon it a few cubic centimetres 
of moderately concentrated sulphuric acid. To this mixture 
add slowly a few cubic centimetres of ordinary alcohol. Notice 
the odor. 

Soap. 

Experiment 245.— In an iron pot boil a few ounces of lard 
w r ith a solution of 40 grams caustic soda in 250 cc. of water for 
an hour or two. After cooling add a strong solution of sodium 
chloride. The soap formed will separate and rise to the top of 
the solution, where it will finally solidify. Dissolve some of the 
soap thus obtained in water. 

Hard Water. 

Experiment 246. — Make some hard water by passing carbon 
dioxide through dilute lime-water until the precipitate first formed 
is dissolved again. Filter. Make a solution of soap by shaking 
up a few shavings of soap with water. Filter. Add the solution 
of soap to the hard water. Is a precipitate formed ? Rub a piece 
of soap between the hands wet with the hard water. 

Experiment 247. — Make some hard water by shaking a litre or 
two of water with a little powdered gypsum. Perform with it the 
same experiments as those first performed with the water con- 
taining calcium carbonate. 



CHAPTER XXXVI. 
OTHER COMPOUNDS OF CARBON. 

The Carbohydrates. — The carbohydrates form an im- 
portant group of carbon compounds which include the 
most abundant substances found in the vegetable kingdom. 
Besides carbon, they generally contain hydrogen and oxy- 
gen in the proportions to form water. Hence they are 
called carbohydrates. The chief compounds included 
under this head are grape-sugar or glucose, cane-sugar, 
starch, cellulose, gum, and dextrin. 

Grape-sugar, Glucose, Dextrose, C 6 H 12 6 . — Dextrose 
occurs very widely distributed in the vegetable kingdom, 
particularly in sweet fruits. It is found also in honey 
and, further, in the liver and the blood. 

Formation of Dextrose. — Dextrose or glucose is formed 
from several of the carbohydrates by boiling with dilute 
mineral acids, or by the action of ferments. Its formation 
from cane-sugar takes place according to this equation, 
equal quantities of dextrose and levulose being formed: 

C u H B O u + H 2 = C 6 H„0 6 + C 6 H„0 6 . 

Cane-sugar. Dextrose. Levulose. 

Its formation from starch is represented by this equation : 
C 6 H 10 O 5 + H,O = C 6 H 12 O 6 . 

Starch. Dextrose. 

Manufacture of Dextrose or Glucose. — Dextrose is pre- 
pared on the large scale from corn-starch in the United 
States, and from potato-starch in Germany. The change 
is usually effected by boiling with dilute sulphuric acid. 
The acid is afterwards removed by treating with chalk, 

667 



668 COLLEGE CHEMISTRY. 

and filtering. [Explain how this removes the acid.] The 
filtered solutions are evaporated either to a syrupy con- 
sistency, and sent into the market under the names 
"glucose," -mixing-syrup," etc. ; or to dryness, the solid 
product being known as "grape-sugar." 

Properties. — Dextrose crystallizes from concentrated 
solutions, and as seen in commercial "granulated grape- 
sugar" looks very much like granulated cane-sugar. It 
is sweet, but not as sweet as cane-sugar. It is estimated 
that the sweetness of dextrose is to that of cane-sugar as 
3: 5. Under the influence of yeast it ferments, yielding 
mainly alcohol and carbon dioxide. Putrid cheese trans- 
forms it into lactic acid, and then into butyric acid. 

Levulose, Fruit-sugar, C 6 H 12 6 .— This form of sugar 
occurs with dextrose in fruits; and is formed by the action 
of dilute acids or ferments on cane-sugar, which breaks 
up according to the equation 

<W>„ + H 2 = 0,H„0. + C 6 H 15 6 . 

Cane-sugar. Dextrose. Levulose. 

As cane-sugar is found in unripe fruits, it is probable 
that the change represented in the equation takes place 
during the process of ripening. 

Cane-sugar, C 12 H 22 O n . — This well-known variety of 
sugar occurs very widely distributed in nature — in sugar- 
cane, sorghum, the Java palm, the sugar-maple, beets, 
madder-root, coffee, walnuts, hazel-nuts, sweet and bitter 
almonds; in the blossoms of many plants, etc., etc. 

Sugar-refining. — Sugar is obtained mainly from the 
sugar-cane and beets. In either case the processes of 
extraction and refining are largely mechanical. When 
sugar-cane is used, this is macerated with water to dissolve 
the sugar. Thus a dark-colored solution is obtained. 
This is evaporated, and then passed through filters of ' 
bone-black by which the color is removed. The clear 
solution is then evaporated in open vessels to some extent; 
and, finally, in large closed vessels called "vacuum-pans," 



SUGAR, ETC. 66 9 

from which the air is partly exhausted, so that the boiling 
takes place at a lower temperature than is required under 
the ordinary pressure of the atmosphere. The mixture 
of crystals and mother - liquors . obtained from the 
"vacuum-pans " is freed from the liquid by being brought 
into the "centrifugals.'* These are funnel-shaped sieves 
which are revolved rapidly, the liquid being thus thrown 
by centrifugal force through the openings of the sieve, 
while the crystals remain behind and are thus nearly dried. 
The final drying is effected by placing the crystals in a 
warm room. 

Molasses. — The mother-liquors obtained from the 
"centrifugals" are further evaporated, and yield lower 
grades of sugar; and, finally, a syrup is obtained which 
does not crystallize. This is molasses. 

Properties of Sugar. — Sugar crystallizes from water in 
large well-formed prisms. When heated to 210° to 220°, 
it loses water, and is converted into a substance called 
caramel, which is colored more or less brown. When 
boiled with dilute acids, cane-sugar is split into equal parts 
of dextrose and levulose. The mixture of the two is called 
invert-sugar. Yeast gradually transforms cane-sugar into 
dextrose and levulose, and these then undergo fermenta- 
tion. Cane-sugar does not ferment. 

Sugar of Milk, Lactose, C 12 H 22 O u -f H 2 0.— This sugar 
occurs in the milk of all mammals. It is obtained in the 
manufacture of cheese. Cow's milk consists of water, 
casein, butter, sugar of milk, and a little inorganic 
material, in about the following proportions : 

Water 87 per cent. 

Casein , 4 " 

Butter 3^ " 

Sugar of milk 4f " 

Mineral matter . , ■ £ ' ' 

100 

Cheese is made by adding rennet to milk, which causes 
the separation of the casein. The sugar of milk remains 



67o COLLEGE CHEMISTRY. 

in solution, is separated by evaporation, and pnrified by 
recrystallization. It has a slightly sweet taste, and is 
much less soluble in water than cane-sugar. 

Souring of Milk. — Sugar of milk ferments under certain 
circumstances, and is transformed mostly into lactic acid. 
The souring of milk is a result of this fermentation. The 
lactic acid formed coagulates the casein; hence the 
thickening. 

Cellulose, C 6 H 10 O 5 . — Cellulose forms, as it were, the 
groundwork of all vegetable tissues. It presents different 
appearances and different properties, according to the 
source from which it is obtained ; but these differences are 
due to substances with which the cellulose is mixed ; and 
when they are removed, the cellulose left behind is the 
same thing, no matter what its source may have been. 
The coarse wood of trees and the tender shoots of the most 
delicate plants consist essentially of cellulose. Cotton- 
wool, hemp, and flax consist almost wholly of cellulose. 

Properties. — Cellulose does not crystallize, and is in- 
soluble in all ordinary solvents. It dissolves in concen- 
trated sulphuric acid. If the solution is diluted and 
boiled, the cellulose is converted into dextrin and dextrose. 
It will thus be seen that rags, paper, and wood, all of 
which consist largely of cellulose, might be used for the 
preparation of dextrose or glucose, and consequently of 
alcohol. 

Gun-cotton, Pyroxylin, Nitrocellulose. — Cellulose has 
some of the properties of alcohols; among them the power 
to form ethereal salts with acids. Thus, when treated 
with nitric acid it forms several nitrates, just as glycerin! 
forms the nitrate known as nitroglycerin (which see). 
The nitrates are explosive, and are used for blasting under 
the name gun-cotton. 

Collodion. — A solution of gun-cotton in a mixture of 
ether and alcohol is known as collodion solution, which is 
much used in photography. When poured upon any 
surface, such as glass, the ether and alcohol rapidly 
evaporate, leaving a thin coating of gun-cotton. 



STARCH. 671 

Celluloid. — Celluloid is an intimate mixture of gun- 
cotton and camphor. As it is plastic at a slightly elevated 
temperature, it can easily be moulded into any desired 
shape. When cooled it hardens. 

Paper. — Paper in its many forms consists mainly of 
cellulose. The essential features in the manufacture of 
paper are, first, the disintegration of the substances used. 
This is effected partly mechanically and partly by boiling 
with caustic soda. Then the resulting mass is converted 
into pulp by means of knives placed on rollers. The pulp, 
with the necessary quantity of water, is then passed 
between rollers. Eags of cotton or linen are chiefly used 
in the manufacture of paper; wood and straw are also 
used. 

Starch, O 6 H 10 O 5 . — Starch is found everywhere in the 
vegetable kingdom in large quantity, particularly in all 
kinds of grain, as maize, wheat, etc. ; in tubers, as the 
potato, arrowroot, etc. ; in fruits, as chestnuts, acorns, 
etc. 

Manufacture of Starch. — In the United States starch is 
manufactured mainly from maize; in Europe, from 
potatoes. The processes made use of are mostly mechan- 
ical. The maize is first treated with warm water; the 
softened grain is then ground between stones, a stream of 
water running constantly into the mill. The thin paste 
which is carried away is brought upon sieves of silk bolt- 
ing-cloth, which are kept in constant motion. The starch 
passes through with the water as a milky fluid. This is 
allowed to settle when the water is drawn off. The starch 
is next treated with water containing a little alkali, the 
object of which is to dissolve gluten, oil, etc. The mix- 
ture is now brought into shallow, long wooden runs, where 
the starch is deposited, the alkaline water running off. 
Finally, the starch is washed with water, and dried at a 
low temperature. 

Properties. — Starch in its usual condition is insoluble 
in water. If ground with cold water it is partly dissolved. 
If heated with water the membranes of the cells of which 



672 COLLEGE CHEMISTRY. 

the starch is composed are broken, and the contents form 
a partial solution. On cooling, it forms a transparent jelly 
called starch-paste. By dilute acids and ferments starch 
is converted into dextrin, maltose, and dextrose. 

Flour. — Wheat flour, which may serve as an example 
of flour in general, contains water, starch with a little 
sugar and gum, gluten, and a small quantity of mineral 
matter. The finest flour contains about 10 per cent of 
gluten and 70 per cent of starch. Gluten is a substance 
that resembles in many respects the white of eggs, or egg- 
albumin. 

Bread-making. — The chemical changes which take place 
in bread-making are of special interest. Bread is made 
by mixing the flour with water and a little yeast. The 
dough thus prepared is put in a warm place for a time, 
when it rises. The rising is a result of fermentation 
caused by the yeast. A part of the starch contained in 
the flour is converted into sugar, and this is then con- 
verted into alcohol and carbon dioxide by fermentation. 
The alcohol passes oft' for the most part, and the carbon 
dioxide in striving to escape from the thick gummy dough 
fills the mass with bubbles of gas, making it light and 
porous. When the loaf is put into the oven the gases 
contained in it expand, making it still lighter; then the 
fermentation is checked by the heat and no further 
chemical change takes place except on the surface, where 
the substances are partly decomposed and converted into 
a dark-colored product, the crust. 

A FEW COMPOUNDS FROM COAL-TAR. 

Aromatic Compounds. — The fact that benzene, 6 H 6 , 
toluene, C 7 H 8 , and other hydrocarbons are obtained from 
coal-tar has already been mentioned (p. 650). These 
hydrocarbons are the starting-points for the preparation 
of a large number of compounds of carbon which are 
commonly called the " aromatic compounds," as many of 
them loave pleasant aromatic odors. 



COAL-TAR COMPOUNDS. 673 

Nitrobenzene, C,.U,X0 2 .— This substance is formed by 
treating benzene with nitric acid: 

C 6 H 6 + 1IN0, = C 6 ILN0 2 + H ! 0. 

It is a yellow liquid with a pleasant odor like that of the 
oil of bitter almonds. It is much used under the name 
artificial oil of bitter almo?ids. 

Aniline, C.H.NH,. — When nitrobenzene is treated with 
a solution from which hydrogen is given off the oxygen is 
extracted and replaced by hydrogen : 

C 6 H 5 N0 2 + 6H = C 6 H 5 NH 8 + 2H 2 0. 

The product is the substance known as aniline. It is a 
colorless liquid. When it together with a similar sub- 
stance, known as toluidine, is treated with mercuric chlo- 
ride, HgCl 2 , or arsenic acid it is converted into the dye 
magenta, which is the substance from which most of the 
aniline dyes are prepared. 

Aniline Dyes. — Of these a large number are in use. 
They are all derivatives of rosaniline, of which magenta is 
a salt. Aniline dyes of great many different colors are 
made, some of them of great beauty. 

Phenol, Carbolic Acid, C 6 H 6 0. — This familiar substance 
is contained in coal-tar, and is extracted from it by treat- 
ing with caustic soda in which the carbolic acid dissolves. 
When pure it crystallizes in beautiful colorless rhombic 
needles. It has a peculiar, penetrating odor, and is 
poisonous. It is much used as a disinfectant. 

_ . .,, . , ' {."CLH.O. — This substance occurs 

Benzoic Aldehyde, j ! b 

in combination with amygdalin, which is found in bitter 

almonds, laurel-leaves, cherry-kernels, etc. Amygdalin 

belongs to the class of compounds known as glucosides, 

which break up into glucose and other substances. 

Amygdalin itself, under the influence of emulsin, which 

occurs with it in the plants, breaks up into oil of bitter 

almonds, hydrocyanic acid, and dextrose : 



674 COLLEGE CHEMISTRY. 

C B H„NO u + 2H,0 = C 7 H 6 + CNH + 2C 6 H 12 6 . 

Arnygdalin Oil of Hydrocy- Glucose 

bitter anic acid 

almonds 

It is prepared from bitter almonds, which yield about 
1.5 to 2 per cent. It is a liquid which has a pleasant odor. 
It is made artificially from coal-tar, and is used in the 
preparation of artificial indigo. 

Benzoic Acid, C T H 6 2 . — Benzoic acid occurs in gum 
benzoin and in the balsams of Peru and Tolu, and is made 
artificially from coal-tar by oxidizing toluene,* C 7 H 8 . 

Balsams and Odoriferous Resins. — The balsams of Peru 
and Tolu are thick fragrant fluids which are obtained from 
certain trees in South America and elsewhere by cutting 
the bark. Benzoin is a similar substance. These as well 
as myrrh, frankincense, and other substances of the kind 
are used for their odors. The odors are intensified when 
the substances are heated. They are largely used as 
incense. 

Gallic Acid, C 7 H 6 5 . — Gallic acid occurs in sumach, in 
Chinese tea, and many other plants. It is formed by boil- 
ing tannin or tannic acid with suphuric acid. It is pre- 
pared from gall-nuts by fermentation of the tannin 
contained in them. It is closely related to tannin or 
tannic acid. 

Tannic Acid, Tannin, C u H 10 O 3 . — This substance occurs 
in gall-nuts, from which it is extracted in large quantities. 
It is soluble in water. Its solution gives a dark blue-black 
color with iron salts. Tannin is used extensively in 
medicine, in dyeing, in the manufacture of leather and of 
ink. 

Tanning. — The process of tanning consists in treating 
hides, from which the hair has been removed, with an 
infusion of hemlock or oak bark, or of sumach-leaves, in 
which there is tannic acid. The acid combines with 
certain parts of the hides, forming insoluble compounds 

* The name toluene comes from the fact that this hydrocarbon was 
first obtained from the balsam of Tolu. 



COAL-TAR COMPOUNDS. 675 

which remain in the pores, converting the hides into 
leather. 

Indigo. — In several plants which grow in the East and 
West Indies, in South America, Egypt, and other warm 
countries, a substance called indican occurs. When this 
is treated with dilute mineral acids or certain ferments, it 
breaks up into indigo-blue and a substance resembling 
glucose. Commerical indigo contains as its principal in- 
gredient indigo-blue. Indigo- blue is now prepared arti- 
ficially by the aid of complicated processes. 

Naphthalene, C 10 H 8 . — This hydrocarbon is contained in 
coal-tar in large quantity. It is a beautiful white crystal- 
lized substance much used in the preparation of dyes and 
for protecting woollen fabrics from moth. 

Anthracene, U H 10 . — Anthracene like naphthalene is 
obtained from coal-tar. Its chief use is in the preparation 
of artificial alizarin. 

Alizarin, C u H 8 4 . — Alizarin is the well-known dye 
obtained from madder-root. For some years it has been 
made artificially from anthracene, and the cultivation of 
madder has been given up. Madder-root was used for 
dyeing "Turkey-red." Artificial alizarin is almost exclu- 
sively used for this purpose at present. 



Glucosides. — Glucosides are substances that occur in 
nature in the vegetable kingdom. They break down into 
sugar and other compounds under the influence of fer- 
ments and dilute acids. Amygdalin has already been 
mentioned. This breaks down into oil of bitter almonds 
and dextrose. Indican, which yields indigo and dextrose, 
is another example. 

Myronic Acid, another glucoside, is found in the form 
of the potassium salt in black mustard- seed. When treated 
with myrosin, which is contained in the aqueous extract of 
white mustard -seed, potassium myronate is converted into 
dextrose and oil of mustard. 



676 COLLEGE CHEMISTRY. 

Alkaloids. — These compounds occur in plants, and are 
frequently those parts of the plants which are most active 
when taken into the animal body. They are hence some- 
times called the active principles of plants, Many of these 
substances are used in medicine. They all contain nitrogen 
and in some respects resemble ammonia. Only a few of 
the more important alkaloids need be mentioned here. 

Quinine. — This valuable alkaloid is obtained from the 
outer bark of certain trees which grow in Peru. The bark 
is known as Peruvian bark. 

Cocaine is found in cocoa-leaves. Its hydrochloric-acid 
salt has come into prominence in medicine, owing to the 
fact that a small quantity of its solution placed upon the 
eye or the gums or injected beneath the skin causes in- 
sensibility to pain. 

Nicotine occurs in tobacco-leaves in combination with 
malic acid. 

Morphine, codeine and narcotine are the principal alka- 
loids found in opium, which is the evaporated sap that 
flows from incisions in the capsules of the white poppy 
before they are ripe. 



INDEX. 



Acetone, 349. 

Acetylene, 510, 652, 655 ; se 
ries, 652. 

Acid, 156. 

Acid, acetic, 349, 650, 658 ; an 
timonic, 333 ; arsenic, 330 
arsenic, experiments with, 343 
arsenious. 330 ; benzoic, 674 
boric, 341 ; bromic, 190 ; car- 
bolic, 673 : butyric, 659 ; car- 
bonic. 366; chloric, 56, 149; 
chloric, experiments with, 151; 
chlorous, 149 ; chlorplatinic, 
647, 648; chlorsulphuric, 241; 
chromic, 598; chromic, experi- 
ments with, 607 ; citric, 661 ; 
cyanic, 380 ; disilicic, 397; di- 
sulphuric, 232 ; dithionic, 253; 
diuranic, 607 ; ferrihydro- 
cyanic, 630; ferrohydrocyanic, 
629; fluosilicic, 392; fluosilicic, 
preparation of, 400 ; formic, 
658 ; fuming sulphuric, 232 ; 
gallic, 674 ; hydrazoic, 279 ; 
hydriodic, 192 ; hydrobromic, 
188 ; hydrobromic, experi- 
ments with, 201 : hydrochloric, 
135 ; hydrochloric, experi- 
ments with, 144 ; hydrochlo- 
ric, formation of, 144 ; hydro- 
chloric, preparation of, 144 ; 
hydrocyanic, 380 ; hydroflu- 
oric, 197 ; hydrofluoric, ex- 
periments with, 204 ; hydro- 
nitric, 279 ; hydrosulphurous, 
234 ; hypobromous, 190 ; hy- 
pochlorous, 148; hyponitrous, 
284 ; hypophosphoric, 328 ; 
hypophosphorous, 329 ; hypo- 
sulphurous, 234 ; iodic, 193 ; 
iodic, experiments with, 204 ; 
lactic, 661 ; malic, 661 ; man- 



ganic, 199, 615 ; metaboric, 
342 ; metachronous, 595 ; met- 
antimonic, 321, 333 ; meta- 
phosphoric, 321, 327 ; metar- 
senic, 321, 330 ; metastannic, 
578, 581 : metavanadic, 339 ; 
molybdic, 603 ; muriatic, 57 ; 
myronic, 675 ; nitric, 57, 279 ; 
nitric, experiments with, 293 ; 
nitric, preparation of, 293 ; 
nitric, red fuming, 282 ; nitric, 
reduction of, 295 ; nitrohy- 
drochloric, 283 ; nitrosylsul- 
phuric, 225 ; nitrous, 283 ; 
nitrous, experiments with, 
296 ; Nordhausen sulphuric, 
232; oleic, 662; orthoanti- 
monic, 321, 333 ; orthophos 
phoric, 321, 323 ; orthovana 
die, 339 ; osmic, 643 ; oxalic 
660 ; palmitic, 659 ; pentathi 
onic, 235; perbromic, 190 
perchloric, 149 ; perchloric 
preparation of, 152 ; periodic 
173, 195 ; pennant- anic, 199 
616, 618 ; persulphuric, 235 
phosphoric, 323 ; phosphoric 
experiments with, 343 ; phos 
phoric, glacial, 328 ; phos 
phoric, "insoluble," 505 
phosphoric, " reverted," 505 
phosphoric, "soluble," 505 
phosphomolybdic, 604 ; phos 
phorous, 328 ; propionic, 659 
prussic, 380 ; pyroantimonic 
321, 333 ; pyroarsenic, 321 
330 ; pyroligneous, 349, 650 
pyrophosphoric, 321, 327; py 
rosulphuric. 232; selenic, 242 
selenious, 242 ; silicic, 394 
silicic, experiments with, 401 
silicic, insoluble, 395 ; silicic, 

677 



678 



INDEX. 



normal, 394 ; stannic, 580 ; 
stearic, 659 ; sulphantimoni- 
ous, 336 ; sulpharsenious, 
332 ; sulpkocyanic, 382 ; sul- 
phuric, 57, 222 ; sulphuric, 
experiments with, 246 ; sul- 
phuric, solid, 232 ; sulphu- 
rous, 233 ; sulphurous, experi- 
ments with, 247, 248 ; sulphy- 
dric, 211; tannic, 674; tar- 
taric, 661 ; telluric, 2*3 ; tellu- 
rious, 242 ; tetraboric, 342 ; 
tetrahydroxylsulphuric, 231; 
tetrathionic, 235 ; thiosulphu- 
ric, 234 ; triazoic, 279 ; trithi- 
onic, 235 ; tungstic, 605 ; va- 
nadic, 339. 

Acids, 88, 153 ; basicity of, 161 ; 
constitution of, 158 ; defini- 
tion of, 420 ; dibasic, 163 • 
fatty, 658 ; monobasic, 162 ;' 
nomenclature of, 166 ; penta- 
basic, 163 ; poly silicic, 396 ; 
polystannic, 581 ; tetrabasic, 
163 ; tribasic, 163 ; trisilicic, 
397. 

Acids of antimony and arsenic, 
constitution of, 321. 

Acids of phosphorus, constitu- 
tion of, 321. 

Acids of sulphur, constitution 
of, 223. 

Acidum phosphoricum glaciate, 
328. 

Affinity, chemical, 9 ; coeffi- 
cients of, 411. 

Agate, 393. 

Air, 253 ; analysis of, 254, 263 ; 
and life, 259 ; liquid, 261. 

Alabandite, 614. 

Alabaster, 503. 

Albite, 470, 531. 

Alcohol, 655 ; ethyl, 655 ; me- 
thyl, 650, 655. 

Alcohols, 664. 

Aldehyde, acetic, 657 ; formic, 
658. 

Aldehydes, 657. 

Algaroth, powder of, 337. 

Alizarin, 675. 

Alkalies, 153, 440. 

Alkali metals, 456. 

Alkaline earths, 492. 

Alkaloids, 676. 

Allanite, 590. 



Allotropy, 120. 

Allylene, 652. 

Alum, 457, 530 ; burnt, 530 ; 
shale, 530. 

Aluminates, 525. 

Aluminium, 522 ; bronze, 524 ; 
chloride, 524 ; preparation of, 
536 ; hydroxide, 525 ; oxide, 
528 ; silicates, 531 ; sulphate, 
528. 

Alums, 529. 

Alunite, 530. 

Amalgamation process for silver, 
548. 

Amalgams, 567. 

Amethyst, 593. 

Ammonia, 267, 270; composition 
of, 274; determination of com- 
position of, 292 ; experiments 
with, 290; preparation of, 290. 

Ammoniacal liquor, 270, 353. 

Ammonia process for soda, 478. 

Ammonium alum, 53u ; amal- 
gam, 277, 568 ; chloride, 483 ; 
diuranate, 607; hydrosulphide, 
485 ; hydroxide, 274 ; nitrate, 
485 ; magnesium phosphate, 
326, 518; phospho-molybdate, 
604 ; salts, 481 ; salts, forma- 
tion of, 291 ; sulphide, 483. 

Ammonium sulphide group, 215, 
484. 

Ampere's hypothesis, 102. 

Anatase, 390, 398. 

Anhydrides, 194. 

Anhydrite, 503. 

Aniline, 673 ; dyes, 673. 

Anions, 416. 

Annealing of glass, 508. 

Anthracene, 650, 675. 

Antimony, 306 ; and its com- 
pounds, experiments with. 
317 ; blende, 335 ; oxycl 1 >- 
rides, 336, ,844; pentachloride, 
313; pentasulphide. 336; pent- 
oxide, 335 ; salts of, 334 ; sul- 
phides, experiments with, 344: 
tetroxide, 335 ; trichloride, 
312; trioxide, 333; trisulphide, 
335. 

Antimonyl salts, 334; sulphate, 
334. 

Apatite, 300, 494, 504. 

Aqua regia. 283. 

Aragonite, 501. 



INDEX. 



679 



Arbor Satumi, 584. 
Argentan, 638. 

Argentic chloride, 550 ; nitrate, 
550. 

Argentous chloride, 550. 

Argon, 261. 

Aromatic compounds, 672. 

Arrhenius, 422. 

Arsenic, 305; and its compounds, 
experiments with, 316 : disul- 
phide, 323, 332 ; pentasul- 
phide, 323, 333 ; pentoxide, 
332 ; sulphides, experiments 
with, 344 ; trichloride, 312 ■ 
trioxide, 330 ; trioxide, experi- 
ments with, 344 ; trisulphide, 
323, 332. 

Arsenical pyrites, 634. 

Arsenides, 305. 

Arseniuretted hydrogen, 306. 

Arsenopja-ite, 634. 

Arsine, 306, 316. 

Atomic theory, 96, 402 ; 
weights, 17, 19, 105 ; weights 
and specific heat, 423. 

Atoms, 102, 403. 

Aurates, 555. 

Auric chloride, 557; compounds, 
554 ; hydroxide, 554 ; oxide, 
554. 

Aurous chloride, 557 ; com- 
pounds, 554 ; oxide, 554 ; sul- 
phide, 557. 

Avogadro's hypothesis, 102. 

Azote, 250. 

Babbitt's metal, 585. 

Balsams, 674. 

Banca tin, 578. 

Barff's process, 626. 

Barite, 511. 

Barium, 511 ; carbonate, 513.; 
chloride, 511 ; chromate, 602, 
dioxide, 512 ; hydroxide, 511 ; 
oxide, 511 ; peroxide, 512 ; 
phosphates, 513 ; sulphate, 
513 ; sulphide, 513. 

Base, 157. 

Bases, 88, 153 ; acidity of, 163 ; 
constitution of, 158 ; definition 
of, 420 ; diacid, 163 ; monacid, 
163 ; nomenclature of, 167 ; 
triacid, 163. 

Basic -lining process, 625. 

Battery, storage, 587. 



Bauxite, 522. 

Bell-metal, 540. 

Bengal fire, 511. 

Benzene, 356, 652, 672 ; series, 
652. 

Benzine, 651. 

Benzoin, 674. 

Berthollet, 126. 

Beryl, 515. 

Beryllium, see Glucinum. 

Bessemer process, 625. 

Bismuth, 309 ; basic nitrates of, 
337, 344 ; dichloride, 313 ; di- 
oxide, 338 ; experiment with, 
317; nitrate, 310; oxy chloride, 
338 ; pentoxide, 338 ; sub- 
nitrate, 338 ; sulphate, 310 ; 
trichloride, 313 ; trioxide, 337; 
trisulphide, 338. 

Bismuthyl salts, 337. 

Blast furnace, 622. ( 

Black ash, 477. 

Black lead, 348. 

Bleaching by chlorine, 130 ; by 
sulphur dioxide, 237 ; powder, 
498. 

Block tin, 578. 

Blow-pipe, 376 ; experiments 
with, 384. 

Bluestone, 543. 

Blue vitriol, 543, 544. 

Bog iron ore, 631 

Bone-black, 351 ; filters, 352, 
357 ; oil, 650. 

Bones, distillation of, 650. 

Boracite, 341, 519. 

Borax, 339, 342, 480. 

Boron, 339 ; experiments with, 
344 ; nitride, 343 ; phosphate, 
342 ; salts of, 342 ; trichloride, 
340. 

Boryl potassium tartrate, 342. 

Boussingault, 255. 

Boyle's law, 44. 

Brass, 540, 561. 

Braunite, 610, 612. 

Bread-making, 672. 

Bricks, 534. 

Brimstone, crude, 207 ; roll, 
207. 

Britannia metal, 307, 579. 

Bromides, experiments with, 
452. 

Bromine, 187 ; preparation of, 
201 ; water, 188. 



6So 



INDEX. 



Bronze, 540. 

Brookite, 390, 398. 

Brown iron ore, 622. 

Bunsen, 256, 257, 469. 

Bunsen burner, 378. 

Burning, 36. 

Butane, 652. 

Butter, 663 ; of antimony, 313. 

Butylene, 652. 

Butyrum Antimonii, 313. 

Cadmium, 564 ; carbonate, 565 ; 
chloride, 565 ; sulphate, 565 ; 
sulphide, 565. 

Csesium, 469, 470. 

Calcium, 494 ; carbide, 510 
carbonate, 501 ; chloride, 495 
chloride, preparation of, 520 
fluoride, 496 ; hydroxide, 496 
nitride, 509 ; oxide, 496; phos- 
phates, 504 ; silicate, 506 ; sul- 
phate, 503 ; sulphide, 509. 

Calc-spar, 501. 

Calomel, 568. 

Calorie, 40. 

Calorimeter, 39. 

Candle, standard, 373. 

Cane-sugar, 668. 

Caramel, 669. 

Carbohydrates, 667. 

Carbon, 345 ; amorphous, 348 ; 
chemical conduct of, 355 ; di- 
oxide, 359 ; dioxide, experi- 
ments with, 369 ; dioxide, re- 
lations to life, 364;disulphide, 
381 ; experiments on reducing 
power of, 358; monoxide, 367; 
experiments with, 371 ; sili- 
cide, 398. 

Carbonates, 366, 449 ; acid, 366 ; 
basic, 367 ; experiments, 455. 

Carbonyl chloride, 369. 

Carbon silicide, 398. 

Carborundum, 398. 

Carnallite, 457, 459, 515. 

Carnelian, 393. 

Caro's reagent, 235. 

Cassiterite, 577. 

Cast iron, 623; gray, 624; white, 
623. 

Cations, 416. 

Caustic soda, 473. 

Cavendish, 54. 

Celestite, 510, 511. 

Celluloid, 671. 



Cellulose, 667, 670. 

Cements, 509 ; hydraulic, 509. 

Cerite, 590. 

Cerium, 390, 590, 591. 

Cerussite, 583. 

Chalcedony, 393. 

Chalcocite, 538, 545. 

Chalk, 494, 501, 502. 

Chance process for recovery of 
sulphur, 478. 

Charcoal, 348 ; absorption of 
gases by, 351, 357; animal, 
351 ; filters, 352, 357 ; kiln, 
349. 

Charring process, 348. 

Chemical action, 7. 

Chemical analysis, 648. 

Chemical change, 2 ; caused by 
heat, 26. 

Chemical reactions, cause of, 409; 
kinds of, 405 ; ideal, 409. 

Chemistry, 3 ; organic, 346, 649. 

Chili saltpetre, 191, 279, 280, 
473. 

Chloral, 675 ; hydrate, 657. 

Chlorates, 447. 

Chloraurates, 555. 

Chlorides, 130, 432; double, 437; 
experiments with, 452; general 
properties of, 436. 

Chlorination process, 556. 

Chlorine, 126, 142; dioxide, 150; 
experiments with, 142 ; prepa- 
ration of, 127; heptoxide, 151; 
hydrate, 133, 143; liquid, 133; 
monoxide, 150. 

Chloro-acids, 437. 

Chloroform, 654. 

Chlorostannates, 580. 

Chloroform, 654. 

Chlorophyll, 622. 

Chlorplatinates, 647. 

Choke-damp, 364. 

Chromates, 598 ; experiments 
with, 607. 

Chrome alums, 598. 

Chrome yellow, 602. 

Chromic chloride, 596 ; com- 
pounds, 594 ; hydroxide, 596 ; 
iron, 595 ; oxide, 597 ; sul- 
phate, 597. 

Chromite, 527, 595. 

Chromium, 244, 594; trioxide, 
600. 

Chromous compounds, 594. 



INDEX. 



68 1 



Chrysoberyl, 526. 

Cinnabar, 567, 572. 

Clay, 388, 522, 532. 

Uleveite, 536. 

Coal, 352 ; anthracite, 352 ; bitu- 
minous, 352; gas, 372, 383; dis- 
tillation of, 649 ; tar, 353, 649. 

Cobalt, 635 ; sulphide, 636. 

Cobaltic hydroxide, 636 ; oxide, 
636. 

Cobaltite, 636, 637. 

Cobaltous chloride, 636 ; cya- 
nide, 637 ; hydroxide, 636. 

Cocaine, 676. 

Codeine, 676. 

Coke, 350. 

Collodion, 670. 

Columbite, 339. 

Columbium, 339. 

Combination, 22 ; direct, 405. 

Combining weights, 15. 

Combustion, 37 ; experiments 
on, 52. 

Compound, chemical, 28. 

Condy's liquid, 618. 

Constitution, 108, 404. 

Copper, 536 ; acetate, 658 ; basic 
carbonate, 545 ; metallurgy 
of, 538 ; native, 538 : plating, 
545 ; ruby, 538, 542. 

Copperas, 633. 

Coral, 494. 

Corrosive sublimate, 568. 

Corundum, 528. 

Cotunnite, 585. 

Crafts, 524. 

Crocoisite, 583, 595. 

Cryolite, 470, 522. 

Cupric arsenite, 544 ; carbonates, 
545 ; chloride, 541 ; hydrox- 
ide, 542 ; nitrate, 544 ; oxide, 
542 ; sulphate, 543; sulphide, 
545. 

Cuprite, 542. 

Cuprous chloride, 541 ; chloride, 
preparation of, 558 ; hydrox- 
ide, 542; oxide, 541; sulphide, 
545. 

Cyanides, 379. 

Cvanogen, 379 ; experiments 
with, 386. 

Dalton, 96. 
Dalton's law, 44. 
Datholite, 341. 



Davy, 126. 

Deacon's process, 127. 

Decay, 458. 

Decomposition, 23 ; direct, 407 ; 
double, 23, 408. 

Decrepitation, 472. 

De la Bastie glass, 508. 

Deliquescence, 84, 95. 

Developers in photography, 552. 

Devil le, 524. 

Dextrin, 667. 

Dextrose, 667. 

Dialyser, 395. 

Dialysis, 395. 

Diamond, 347. 

Diaspore, 525. 

Didymium, 339, 591. 

Diffusion, 58, 260 ; experiments 
on, 68. 

Dimorphism, 209. 

Dissociation, 87, 311, 414 ; elec- 
trolytic, 417; of dissolved sub- 
stances, 417. 

Distillation, 95 ; destructive, 
349 ; dry, 349. 

Dolomite, 359, 515. 

Double chloride, 437. 

Drummond light, 78. 

Dulong, 426. 

Dumas, 255 ; his study of com- 
bustion of hydrogen, 72 ; 
method for determining the 
specific gravity of vapors, 112. 

Dynamite, 663. 

Earthenware, 534. 

Efflorescence, 84, 95, 474. 

Eka-aluminium, 574. 

Eka-boron, 535. 

Eka- silicon, 577. 

Electrolysis, 416. 

Electrolytes, 416. 

Electrolytic dissociation, 417. 

Electrolytic process for chlorine, 
129. 

Electro-negative ions, 417. 

Electro-positive ions, 417. 

Electrotypes, 546. 

Elements, 6, 17 ; acid-forming, 
141 ; base-forming, 158 ; re- 
placing power, 111 ; symbols 
of, 19. 

Emerald, 515. 

Emery, 528. 

Emulsin, 673. 



682 



INDEX. 



Energy, chemical, 41 ; conserva- 
tion of, 5; stored up in plants, 
865. 

Enstatite, 397. 

Epsom salt, 515, 517. 

Equations, chemical, 22. 

Equilibrium, 412. 

Erbium, 520. 

Etching on glass, 198, 202. 

Ethane, 652. 

Ether, 661. 

Ethereal salts, 662. 

Ethers, 661. 

Bthiops martialis, 572. 

Ethylene, 654 ; series, 652. 

Eudiometer, 73. 

Eudiometric method, 73 ; ex- 
periments, 81. 

Faraday's law, 416. 

Fats, 662. 

Feldspar, 451, 457, 522, 531. 

Ferment, nitrifying, 280. 

Fermentation, 650, 656 ; acetic 
acid, 656 ; alcoholic, 656 ; lac- 
tic acid, 656. 

Ferric chloride, 627 ; chloride, 
preparation of, 640 ; ferro- 
cyanide, 630 ; hydroxide, 630 ; 
soluble, 631 ; oxide, 632 ; sul- 
phate, 633 ; sulphate, prepara- 
tion of, 640. 

Ferrous ammonium sulphate, 
preparation of, 640 ; carbon- 
ate, 632 ; chloride, 627 ; ferric 
oxide, 631 ; ferricyanide, 630 ; 
hydroxide, 630 ; phosphate, 
633 ; sulphate, 633 ; sulphate, 
preparation of, 640 ; sulphide, 
632. 

Fire-damp, 653. 

Flame, 373 ; oxidizing, 375 ; re- 
actions, 487 ; reducing, 375. 

Flames, 373 ; causes of lumin- 
osity of, 377 ; structure of, 
375. 

Flint, 393. 

Flores zinci, 564. 

Flour, 672. 

Fluorine, 196. 

Fluor-spar, 494, 496. 

Flux, 197, 496, 502. 

Fool's gold, 634. 

Formulas, constitutional, 109 ; 
molecular, 107, 403. 



Frankincense, 674. 
Franklinite, 560, 631. 
Fulminating mercury, 572 . 
Fumaroles, 341. 
Fusel oil, 656. 

Gadolinite, 590. 

Gahnite, 526, 560. 

Galena, 583. 

Galenite, 583. 

Galvanized iron, 561. 

Gallium, 574. 

Garnet, 506 

Gas, normal, 104. 

Gases, kindling temperature of, 
374 ; measurement of volume 
of, 43. 

Gasoline, 651. 

Gay Lussac, 126 ; law of, 44 ; 
tower, 227. 

Germanium, 577 ; chloride, 577. 

German silver, 541, 638. 

Gersdorffite, 638. 

Glass, 506. 

Glauber, 135, 474. 

Glauber's salt, 474. 

Glover tower, 227. 

Glucinum, 515. 

Glucose, 667. 

Glucosides, 675. 

Gluten, 672. 

Glycerin, 657. 

Gold, 554 ; alloys, 557 ; amal- 
gam, 557 ; metallurgy of, 555 ; 
sulphide, 557. 

Goldschmidt's welding process, 
523. 

Graham, 61. 

Granite, 531. 

Grape-sugar, 667. 

Graphite, 347. 

Greenockite, 565. 

Green vitriol, 544, 68a 

Gu 11 berg, 411. 

Gum, 667. 

Gun-cotton, 670. 

Gun-metal, 540. 

Gunpowder, 463. 

Gypsum, 448, 494, 508. 

Halogens, 186. 

Hardness, permanent, 502, 504 ■ 

temporary, 502. 
Hard water, 502 ; experiments 

with, 666. 



INDEX. 



6*3 



Hauerite, 614. 

Hausmannite, 610, 612. 

Heat, of combustion. 39 ; of de- 
composition, 40; of neutraliza- 
tion, 418 ; specific, law of, 
4-25. 

Heavy spar, 448, 511. 

Helium, 536. 

Hematite, (522, 632. 

Heptane, 652. 

Hexane, 052. 

Homologous series, 052. 

Homology, 652. 

Hornblende, 515. 

Hydrargilite, 525. 

Hydrates, 87, 167, 169. 

Hydraulic cement, 509 ; mining, 
555. 

Hydrazine, 278. 

Hydrazoic acid, 279. 

Hydrocarbons, 356, 358, 050; 
formation of, 049. 

Hydrochloric acid hydrate, 138. 

Hydrogen, 54 ; amount evolved 
when a known weight of a 
metal is dissolved in an acid, 05; 
burning of, 72; chemical prop- 
erties of, 59, 09 ; dioxide, 122; 
dioxide, experiments with, 
124 ; peroxide, 122 ; persul- 
phide, 217; preparation of, 54, 
02 ; sulphide, 211 ; sulphide, 
experiments with, 221 ; sul- 
phide group, 214 ; sulphide in 
chemical analysis, 214. 

Hydrogenium, 01. 

Hydrogen- valence, 182. 

Hydrosulphides, 210. 445. 

Hydroxides, 440 ; experiments 
with, 453 ; formation from 
oxides, 239. 

Hydroxyl, 004. 

Hydroxylainine, 278. 

Hypochlorites, 447. 

Hyposulphite of soda, 475. 

Iceland spar, 501. 
Ice-machine, 271. 
Illuminating gas, 372. 
Illumination, 372. 
Incense, 074. 
Indican, 075. 
Indigo, 075 
Indium, 575. 
Infusorial earth, 394. 



Ink, sympathetic, 636. 

Invert- sugar, 009. 

Iodic anhydride, 194. 

Iodides, experiments with, 452. 

Iodine, 190; bromide, 190; chlo- 
ride, 190 ; experiments with, 
202 ; pentoxide, 194 ; trichlo- 
ride, 190. 

Iodoform, 054. 

Ions, 90, 410, 417. 

Iridium, 644. 

Iron, 621 ; disulphide, 634 ; gal- 
vanized, 501; metallurgy of, 
622 ; ores, 622 ; pyrites, 634. 

Kainite, 400, 515. 

Kaolin, 522, 532. 

Kelp, 190. 

Kieserite, 515, 517. 

Kindling temperature, 37 ; of 

gases, 374. 
King's yellow, 332. 
Kirchhoff, 409. 
Kruegel, 262. 
Krypton, 262. 

Lactose, 669. 

Ladenburg, 262. 

Lamp-black, 350. 

Lanthanum, 536, 590, 591. 

Lapis lazuli, 533. 

Laughing gas, 285. 

Lavoisier, 32, 250, 254. 

Law of definite proportions, 12, 
402 ; of Dulong and Petit, 
425 ; of indestructibility of 
matter, 4, 402 ; of multiple 
proportions, 13, 402 ; of speci- 
fic heats, 425. 

Lead, 582 ; acetate, 058; carbon- 
ate, 588 ; chloride, 585, 592 ; 
chromate, 002 ; hydroxide, 
580 ; iodide, 585, 592 ; metal- 
lurgy of, 583; molybdate, 003 ■ 
nitrate, 588 ; oxide, 585 ; pen- 
cils, 348 ; peroxide, 580, 592 ; 
sesquioxide, 592 ; suboxide, 
585 ; sugar of, 058 ; sulphate, 
589 ; sulphide, 588 ; tetrachlo- 
ride, 585 ; tree, 584. 

Le Blanc process for soda, 470. 

Lepidolite, 409, 481. 

Levulose, 008. 

Lewies, 378. 



6S 4 



INDEX. 



Lignite, 352. 

Lime, 496 ; air-slaked, 496 ; 

chloride of, 498 ; slaked, 496. 
Lime- kiln, 496. 
Lime-light, 78. 
Limestone, 494, 501. 
Lime water, 497. 
Linking of atoms, 404. 
Linngeite, 637 
L quor ferri chlorati, 627 ; ferri 

sesquichlorati, 628. 
Litharge, 586 
Lithium, 481 ; carbonate, 481 ; 

chloride, 481 ; phosphate, 

481 
Loadstone, 631 
Lunar caustic, 553. 

Magenta, 673. 

Magnalium, 534. 

Magnesia, 517 ; alba, 518 ; usta, 
517. 

Magnesio-ferrite, 527. 

Magnesite, 515. 

Magnesium, 515 ; borates, 519 
carbonate, 518 ; chloride, 516 
chloride, preparation of, 520 
hydroxide, 517 ; oxide, 517 
phosphates, 518; silicates, 519 
silicide, 519 ; sulphate, 517. 

Magnetite, 527, 622, 631. 

Malachite, 538, 545. 

Malaria, 261. 

Manganates, 615. 

Manganese, 199, 609 ; black 
oxide of, 199, 610, 612 ; di- 
oxide, 199, 610, 612; disul- 
phide, 614 ; heptoxide, 201, 

618 ; salts, preparation of, 

619 ; tetrachloride, 611 ; tetra- 
chloride, 611 ; trichloride, 611. 

Manganic oxide, 612 ; sulphate, 

615. 
Manganite, 610, 612. 
Manganous chloride, 610, 619 ; 

hydroxide, 612 ; manganic 

oxide, 612 ; sulphate, 614 ; 

sulphide, 614. 
Marble, 494, 501. 
Marcasite, 634. 
Marl, 501, 532. 
Marsh- gas, 356, 653. 
Marsh's test for arsenic, 308, 

317. 
Mass action, 411. 



Mass, influence of, 410. 

Matches, safety, 303. 

Matter, constitution of, 5 ; law 
of indestructibility of, 4. 

Meerschaum, 515. 

Mendeleeff, 173. 

Mercuric chloride, 569 ; com- 
pounds, 566 ; cyanide, 572 ; 
diammonium chloride, 573 ; 
iodide, 570 ; nitrate, 573 ; ni- 
trate, preparation of, 575 ; ox- 
ide, 571 ; sulphide, 572. 

Mercurius solubilis Haline- 
manni, 573. 

Mercurous ammonium chloride, 
573 ; chloramide, 573 ; chlo- 
ride, 568 ; compounds, 566 ; 
iodide, 570 ; nitrate, 573 ; ni- 
trate, preparation of, 575. 

Mercury, 566 ; metallurgy of, 
567. 

Meta- arsenates, 330. 

Metallic properties, 428. 

Metallurgy, 430. 

Metals, 158, 428 ; properties, 
431. 

Metaphosphates, 326. 

Metastannates, 581. 

Metathesis, 23, 408. 

Meteorites, 621, 637. 

Methane, 356, 653. 

Methyl alcohol, 349. 

Meyer, Lothar, 173. 

Meyer, Victor, method for de- 
termining the specific gravity 
of vapors, 113. 

Mica, 506, 522, 531. 

Microcosmic salt, 485. 

Milk, souring of, 670. 

Minerals, 430. 

Miner's safety lamp, 374. 

Minium, 587. 

Mixtures, mechanical, 10, 28. 

Moissan, 347. 

Molasses, 669. 

Molecular weights, 103. 

Molecules, 102, 403. 

Molybdates, 603. 

Molybdenite, 602. 

Molybdenum, 244, 602. 

Morley's determination of the 
ratio between the atomic 
weights of hydrogen and oxy- 
gen, 73. 

Morphine, 676. 



INDEX. 



685 



Mortar, 508. 
Mosaic gold, 582. 
Murium, 126. 
Myrrh, 674. 

Naphtha, 651. 

Naphthalene, 650, 675. 

Narcotine, 676. 

Nasceut state, 121. 

Neodyrnium 339, 591. 

Neon, 262. 

Neutralization, 153 ; experiments 
on, 170. 

Newton's metal, 310. 

Niccolite, 638. 

Nickel, 637 ; alloys, 638 ; car- 
bonyl, 639 ; cyanide, 638 ; 
plating, 638. 

Nickelic hydroxide, 638. 

Nickelous hydroxide, 638 ; ox- 
ide, 638. 

Nicotine, 676. 

Nilson, 524. 

Niobium, 339. 

Nitrates, 446 ; formation of, 268 ; 
formation of, in the soil, 268, 
280, 462. 

Nitric oxide, 285 ; experiments 
with, 296. 

Nitrification, 280, 462. 

Nitrobenzene, 673. 

Nitrocellulose, 670. 

Nitrogen, 250 ; boride, 343 ; 
chloride, 289 ; experiments 
with, 262 ; iodide, 290 ; pent- 
oxide, 289 ; peroxide, 287 ; 
peroxide, experiments with, 
297 ; preparation of, 251, 262; 
relations to life, 259 ; trioxide, 
287 ; trioxide, experiments 
with, 297. 

Nitroglycerin, 663. 

Nitrosyl chloride, 283. 

Nitrous oxide, 284, experiments 
with, 296. 

Non-metals, 428. 



Octane, 652. 

Oil of bitter almonds, 673 ; arti- 
ficial, 673. 
Olefiant gas, 356, 654. 
Oleomarearin, 663. 
Olivine, 519. 
Opal, 393. 



Opium, 676. 

Ores, 430. 

Orpiment, 332. 

Orthoclase, 398. 

Osmium, 643 ; peroxide, 643. 

Osmotic pressure, 422. 

Oxidation, slow, 38. 

Oxides, 42, 439 ; acidic, 194 ; 
basic, 194. 

Oxygen, 32 ; amount liberated 
from a known weight of po- 
tassium chlorate, 47 ; burning 
of, 60 ; chemical properties of, 
35, 50 ; physical properties of, 
35, 50 ; preparation of, 32, 
42 ; valence, 182. 

Oxyhydrogen light, 78, 82; blow- 
pipe, 77 ; blow-pipe, experi- 
ments with, 81. 

Ozone, 116 ; in the air, 119. 

Palladium, 644 ; hydrogen, 644. 

Paper, 671. 

Paraffin, 651. 

Paraldehyde, 657. 

Paris green, 544. 

Passive state of iron, 627. 

Pattinson's method of separating 

silver from lead, 548. 
Peat, 352. 
Pentane, 652. 
Periodates, 195. 
Periodic law, 173, 405. 
Permanent white 513. 
Permanganates, 616. 
Petalite, 481. 
Petit, 425. 

Petroleum, 649, 651. 
Fettersson, 524. 
Pfeffer, 422. 
Phenol, 673. 
Philosopher's wool, 563. 
Phosgene, 369. 
Phosphates, 450; acid, 325 

neutral, 325 ; normal, 325 

primary, 325 ; secondary, 325 

tertiary, 325. 
Phosphine, 304 ; experiments 

with, 315 ; liquid, 305. 
Phosphoric anhydride, 329. 
Phosphorite, 300, 494, 504 
Phosphorous anhydride, 329. 
Phosphorus, 300 ; experiments 

with, 314 ; oxychloride, 330 ; 

pentachloride, 311 ; pentachlo- 



686 



INDEX. 



ride, preparation of , 319; pent- 
oxide, 329 ; red, 303 ; trichlo- 
ride, 310 ; trichloride, prep- 
aration of, 317 ; trioxide, 
329. 

Phosphorus, constitution of 
acids of, 321. 

Phosphuretted hydrogen, 304. 

Photography, 552. 

Photometer, 373. 

Physical change, 2. 

Physics, 3. 

Pig-iron, 623. 

Pinchbeck, 540. 

Pink salt, 580. 

Pitch blende, 605. 

Placer mining, 555. 

Plaster of Paris, 503. 

Platinic chloride, 647 ; sulphide, 
648. 

Platinous chloride, 647; sulphide, 
648 

Platinum, 645 ; alloys, 647 ; 
black, 646 ; metals, 642 ; 
spongy, 645. 

Plumbago, 347. 

Plumbates, 587. 

Plumbites, 586. 

Pollux, 469. 

Porcelain, 532, 533. 

Potash, 457, 467. 

Potassium, 457 ; alum, 530 ; bi- 
chromate, 599, 607 ; bromide 
459; carbonate, 467; carbonate 
acid, 468 ; carbonate, extrac 
tion from wood-ashes, 490 
chlorate, 464"; chlorate, prep 
aration of, 151 ; chloride, 459 
chromate, 599, 607 ; cyanide 
465 ; ferricyanide, 630 ; ferro 
cyanide, 379, 629 ; fluoger 
manate, 577 ; fluoride, 459 
fiuosilicate, 487; hydride, 458 
l^drosulphide, 461 ; hydrox 
ide, 460 ; hypochlorite, 147 
iodide, 459 ; iodide, prepara 
tion of, 489 ; manganate, 615 
manga nate, preparation of 
619 ; nitrate, 462 ; nitrite 
463 ; oxide, 461 ; perchlorate 
465 ; perchlorate, preparation 
of, 152 ; permanganate, 200 
617 ; permanganate, prepara 
tion of, 619 ; peroxide, 461 
phosphates, 468 ; silicate, 469 



sulphate, 466 ; sulphate, acid, 
466 ; sulphite, 467 ; sulphite, 
acid, 467 ; sulphocyanate, 
466. 

Praseodymium, 339, 591. 

Preparing salt, 581. 

Priestley, 32, 254. 

Printing ink, 350. 

Propane, 652. 

Propylene, 652. 

Prussian blue, 630. 

Puddling, 624. 

Purple of Cassius, 557.. 

Pyrargyrite, 553. 

Pyrite, 622, 634. 

Pyroarsenates, 330. 

Pyrolusite, 199, 610, 612. 

Pyrophosphates, 327. 

Pyrosiderite, 631. 

Pyroxylin, 670. 

Pyrrhotite, 622. 

Quartation, 556. 
Quartz, 388. 
Quartzite, 388, 393 
Quicklime, 496. 
Quinine, 676. 

Radicals, 664. 

Ramsay, 261, 262. 

Rayleigh, 261. 

Raoult's methods for determin- 
ing molecular weights, 420, 
422. 

Reaction, chemical, 409. 

Realgar, 332. 

Red fire, 511. 

Red lead, 587. 

Red phosphorus, 303. 

Red prussiate of potash, 638. 

Reducing agent, 60. 

Reduction, 60, 70 ; by carbon, 
356. 

Residues, 664. 

Respiration. 363. 

Reversion of phosphates, 505. 

Rhodium, 644. 

Rhodocroisite, 610. 

Rinmann's green, 564. 

Roasting, 445 

Rock-crystal, 393. 

Rose's metal, 310. 

Rouge, 632. 

Rubidium, 469, 470. 



INDEX. 



687 



Ruby, 528. 

Ruby copper, 538, 542. 
Rupert's drops, 508. 
Ruthenium, 643. 
Kutile, 390, 398. 

Safety lamp, 374. 

Safety matches, 303, 465. 

Sal ammoniac, 270, 483. 

Salt, common, 471. 

Saltpetre, 457, 462 ; plantations, 
462. 

Salts, 157 ; acid, 162, 164 ; basic, 
163, 164 ; constitution of, 
160 ; decomposition by bases, 
441 ; formation of, 434 ; neu- 
tral, 162 ; nomenclature of, 
168 ; normal, 162, 164. 

Samarium, 536. 

Samarskite, 536. 

Sand, 388, 393. 

Saponification, 662. 

Sapphire, 528. 

Scandium, 535. 

Scheele, 32, 126, 254. 

Scheele's green, 544. 

Scheelite, 604. 

Schlippe's salt, 336. 

Schweinfurt green, 544. 

Selenium, 218 ; dioxide, 242. 

Serpentine, 515. 

Siderite, 622, 632. 

Siemens-Martin furnace, 625. 

Silica, 388. 

Silicates, 388, 451. 

Silicides, 398. 

Silicon, 388 ; dioxide, 388, 393 ; 
hexachloride, 391 ; hydride, 

390 ; magnesium, 519 ; prepa- 
ration of, 399 ; tetrachloride, 

391 ; tetrafluoride, 392 ; tetra- 
fluoride, preparation of, 400. 

Silver, 546 ; allotropic forms of, 
549 ; alloys of, 550 ; amal- 
gam, 550 ; bromide, 550 ; 
chloride, 550 ; cyanide, 553 ; 
iodide, 551 ; metallurgy of, 
547 ; nitrate, 553 ; oxide, 552 ; 
peroxide, 553 ; suboxide, 553 ; 
tree, preparation of, 558 ; 
triazoate, 552. 

Slaking, 497. 

Smalt, 637. 

Smaltite, 636. 

Smithsonite, 560, 564. 



1 Soap, action on hard waters, 

659 ; preparation, 666. 
I Soaps, 659. 

Soapstone, 515. 

Soda, 475 ; calcined, purified, 
477 ; crude, 477 ; crystallized, 
477. 

Soda-water, 362. 

Sodium, 470 ; alum, 531 ; amal- 
gam, 567 ; ammonium phos- 
phate, 485 ; bichromate. 602 ; 
borate, 480 ; bromide, 473 ; 
carbonate, 475 ; carbonate, 
acid or primary, 479 ; carbo- 
nate, preparation of by am- 
monia process, 478 ; chloride, 
471; chromate, 602; diuranate, 
607 ; fluoride, 473 ; hydride, 
471 ; hydroxide, 473 ; iodide, 
473 ; oxide, 473 ; nitrate, 473 ; 
peroxide, 473 ; phosphates, 
479 ; silicate, 480 ; stannate, 
580 ; sulphantimonate, 473 ; 
sulphate, 473 ; sulphate, ex- 
periment on supersaturated 
solution of, 490 ; thiosulphate, 
475. 

Solder, soft, 579. 

Solution, 88 ; as an aid to chem- 
ical action, 90. 

Solvay process for soda, 478. 

Spathic iron, 632. 

Specific heat and atomicweights, 
423. 

Specific gravity of vapors, de- 
termination of, 112. 

Spectroscope, 487. 

Spectrum, continuous, 488 ; line, 
488. 

Spelter, 560. 

Sphalerite, 560. 

Spiegel iron, 624. 

Spinels, 526. 

Spirits of wine, 655. 

Spiritus fumans Libarii, 580. 

Spitting of silver, 549. 

Spodumene, 481. 

Stalactites, 502. 

Stalagmites, 502. 

Stannic chloride, 580 ; chloride 
preparation of , 592 ; hydrox- 
ide, 580 ; oxide, 581 ; sul- 
phide, 581. 

Stannite, 578. 

Stannous chloride, 579; chloride; 



688 



INDEX. 



preparation of, 592 ; sulphide, 

581. 
Starch, 667, 671. 
Stearin, 659. 
Steel, 625. 
Stibine, 307, 317. 
Stibnite, 306, 335. 
Stolzite, 605. 
Storage battery, 587. 
Strass, 507. 
Stromeyerite, 553. 
Strontianite, 510. 
Strontium, 510 ; chloride, 510 ; 
hydroxide, 510 ; nitrate, 510 ; 
oxide, 510 ; sulphate, 511. 
Sublimation, 483. 
Substitution, 55. 
Sugar of lead, 658; of milk, 669; 

refining, 668. 
Sulphantimonates, 336. 
Sulphates, 448 ; experiments 

with, 454. 
Sulphides, 210, 443. 
Sulphites, 449 ; acid, 233. 
Sulpho-salts, 446. 
Sulphostannates, 582. 
Sulphur, 206 ; acid chlorides of, 
240; auratum/Sift', dichloride, 
218 ; dioxide, 235 ; dioxide, ex- 
periments with, 247 ; flowers 
of, 207; heptoxide, 235 ; liexa- 
tluoride, 218 ; insoluble, 210 ; 
monochloride, 218 ; sesquiox- 
ide, 235 ; soluble, 210 ; stick, 
207 ; tetrachloride, 218 ; triox- 
ide, 238; trioxide, experiments 
with, 248 ; waters, 212. 
Sulphur, constitution of acids of , 

223. 
Sulphuretted hydrogen, 211. 
Sulphuryl chloride, 240. 
Sulphuryl hydroxyl chloride, 

241. 
Sulphydrates, 217. 
Superphosphate of lime, 505. 
Sylvite, 457, 459. 
Symbols of compounds, 20 ; of j 

elements, 19. 
Sympathetic inks, 636. 

Tachydrite, 495. 
Talc, 515. 
Tannin, 674. 
Tanning, 674. 
Tantalite, 339. 



Tantalum, 339. 

Tartar, cream of, 661 ; crude 
457 ; emetic, 334. 

Tellurium, 219 ; dioxide, 244 ; 
monoxide, 244 ; trioxide, 244. 

Tempering of glass, 508; of steel, 
625. 

Thallium, 575. 

Thenard, 126. 

Theory, use and value of a, 99. 

Thomas- Gilchrist process, 625. 

Thorite, 390, 397. 

Thorium, 390. 

Tin, 577 ; amalgam, 579 ; metal- 
lurgy of, 578 ; salt, 579 ; stone, 
577, 581. 

Titanium, 390 ; dioxide, 395. 

Toluene, 652, 672. 

Toluidine, 673. 

Tridymite, 393. 

Troost, 524. 

Tungstates, 605. 

Tungsten, 244, 604. 

Turkey red, 675. 

Turnbull's blue, 630. 

Tuyeres, 622. 

Type metal, 307. 



Ultramarine, 533. 

Uranates, 607. 

Uraninite, 606. 

Uranium, 244, 605 ; oxides, 606 ; 

yellow, 607. 
Uranous salts, 606. 
Uranyl nitrate, 606 ; sulphate, 

607 ; salts, 606. 

Valence, 109, 404 ; variations of, 

404. 
Vanadium, 338. 
Van't Hoff, 423. 
Vein-mining, 555. 
Ventilation, 260. 
Verdigris, 658. 
Vinegar, 658 ; mother of, 656, 

658. 
Vitriol, oil of, 57, 223. 
Vitriols, 544. 
Vivianite, 633. 

Waage, 411. 

Water, 83 ; as a solvent, 88 ; con- 
tamination of, by sewage, 92 ; 



MAY 261949 



INDEX. 



689 



determination of composition 
of, 79; in organic substances, 

83, 94; of crystallization, 83, 
94 ; proofs of composition of, 
84 ; purification of, 93, 95 ; 
synthesis of, 85. 

Water-gas, 368. 

Water-glass, 469, 481. 

Waters, chalybeate, 92 ; effer- 
vescent, 92 ; natural, 91 ; sul- 
phur, 92. 

Weights, atomic, 17, 100 ; com- 
bining, 15, 100. 

Weldon's process, 128, 613. 

Welsbach, Auer von, 591. 

White arsenic, 330. 

White lead, 588. 

White precipitate, 573 ; fusible, 
573 ; infusible, 573. 

White vitriol, 544, 564. 

Winkler, 576. 

Witherite, 511. 

Wolframite, 604. 

Wollastonite, 397, 506. 

Wood, distillation of, 650. 

Wood's metal, 310. 



Wood spirit, 349, 650. 655. 
Work, chemical, 41. 
Wrought iron, 624, 625. 
Wulfenite, 583, 603. 

Xenon, 262. 
Xylene, 652. 

Yellow prussiate of potash, 379, 

628. 
Ytterbium, 536. 
Yttrium, 536. 

Zeolites, 506. 

Zinc, 560 ; alloys, 561 ; amal- 
gam, 561 ; blende, 563 ; car- 
bonate, 564 ; chloride, 562 ; 
dust, 560 ; experiment on 
burning of, 575 ; hydroxide, 
562 ; metallurgy of, 560 ; ox- 
ide, 563 ; solution of, in so- 
dium hydroxide, 575 ; sul 
phate, 564 ; sulphide, 563 ; 
white, 563. 

Zircon, 390, 397. 

Zirconium, 390. 



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